7 


METHUEN'S  TEXT-BOOKS    OF    SCIENCE 


A    TEX  T-B  O  O  K    OF 
INORGANIC    CHEMISTRY 


TEXT-BOOKS  OF  SCIENCE 

Crown  8vo. 

THE   COMPLETE  SCHOOL  CHEMISTRY.     By  F.  M.  OLDHAM,  B.A.     Seventh 

Edition,  Revised.    4*.  f>d. 
A    TEXTBOOK  OF  INORGANIC   CHEMISTRY.     By  G.  SENTER,  D.Sc,,  Ph.D. 

Fifth  Edition.     75.  6d. 

PRACTICAL  CHEMISTRY.     Part  I.     By  W.  FRENCH,  M.A.    Sixth  Ed.     is.  6d. 
PRACTICAL  CHEMISTRY.     Part  II.    By  W.  FKENCH,  M.A.,  and  T.  H.  BOARD- 
MAN,  M.A.     is.  6d. 
A  TEXTBOOK  OF  PRACTICAL  CHEMISTRY  FOR  TECHNICAL  INSTITUTES. 

By  A.  E.  DUNSTAN,  D.Sc.,  and  F.  B.  T.  THOLE,  B.Sc.     3s.  6d. 
AN  INTRODUCTION  TO  QUANTITATIVE  ANALYSIS.     By  S.  J.  M.  AULD, 

D.Sc.,  Ph.D.     5s. 
A  FIRST  YEAR  COURSE   OF    ORGANIC    CHEMISTRY   FOR    TECHNICAL 

INSTITUTES.     By  A.  E.  DUNSTAN,  D.Sc.     2s.  6d. 
A  SECOND  YEAR  COURSE  OF  ORGANIC  CHEMISTRY  FOR  TECHNICAL 

INSTITUTES.     By  F.  B.  T.  THOLE,  B.Sc.     23.  6d. 
A  THIRD  YEAR   COURSE  OF  ORGANIC  CHEMISTRY  FOR  TECHNICAL 

INSTITUTES.     By  T.  P.  HILDITCH,  D.Sc.     3s. 

QUALITATIVE  ORGANIC  ANALYSIS.     By  F.  B.  T.  THOLE,  B.Sc.     is.  6d. 
MODERN  RESEARCH  IN  ORGANIC  CHEMISTRY.   By  F.  G.  POPE,  B.Sc.,  F.C.S. 

75.  6d. 

A  FIRST  YEAR  PHYSICAL  CHEMISTRY.     By  T.  P.  HILDITCH,  D.Sc.    as. 
OUTLINES  OF  PHYSICAL  CHEMISTRY.      By  GEORGE  SENTER,  B.Sc.  (Lond.), 

Ph.D.    Ph.D.     Sixth  Edition,  Revised.     6s. 

PHYSICO-CHEMICAL  CALCULATIONS.     By  JOSEPH  KNOX,  D.Sc.     as.  6d. 
ELEMENTARY  SCIENCE  FOR  PUPIL  TEACHERS.     Physics  section  by  W.  T. 

CLOUGH,  A.R.C.S.    Chemistry  section  by  A.  E.  DUNSTAN,  D.Sc.    as. 
FIRST  YEAR  PHYSICS.     By  C.  E.  JACKSON,  M.A.     Second  Edition,     is.  6d. 
EXAMPLES  IN  PHYSICS.     By  C.  E.  JACKSON,  M.A.     Third  Edition.     2j.  6d. 
EXAMPLES   IN   ELEMENTARY  MECHANICS:  Practical,  Graphical,   and 

Theoretical.    By  W.  J.  DOBBS,  M.A.    55. 

PRACTICAL  MECHANICS.     By  SIDNEY  H.  WELLS.     Sixth  Edition.     35.  6d. 
A  HANDBOOK  OF  PHYSICS.     By  W.  H.  WHITE,  M.A.    7s.  6d. 
PRELIMINARY  PHYSIOLOGY.     By  W.  NARRAMORE,  F.L.S.    3s.  6d. 
PLANT  LIFE :  Studies  in  Garden  and  School.  By  HORACE  F.  JONES.  F.C.S.  35. 6d. 
A  CONCISE  HISTORY  OF  CHEMISTRY.     By  T.  P.  HILDITCH,  D.Sc.     25.  6d. 
ELEMENTARY  CHEMICAL  THEORY.     By  J.  M.  WADMORE,  M.A.     35.  6d. 
TECHNICAL  ARITHMETIC  AND  GEOMETRY.     By  C.  T.  MILLIS,  M.I.M.E. 

Second  Edition.     35.  6d. 

TEXT-BOOKS  OF  TECHNOLOGY 

ELECTRIC  LIGHT  AND  POWER.      By  E.    E.  BROOKS,  B.Sc.,   Head  of  the 

Department  of  Physics  and  Electrical  Engineering,  Leicester  Technical  School, 

and  W.  H.  N.  JAMES,  A.R.C.S.,  Municipal  School  of  Technology,  Manchester. 

With  17  plates  and  210  diagrams.     Third  Edition.     Crown  8vo.    45.  6d. 
ENGINEERING    WORKSHOP  PRACTICE.!    By  C.   C.    ALLEN,   Head  of   the 

Engineering  Department,  Technical  Institute,  Auckland.  With  152  illustrations. 

Second  Edition.     Crown  8vo.     35.  6d. 
CARPENTRY  AND  JOINERY.    By  F.  C.  WEBBER,  Chief  Lecturer  to  the  Building 

Trades  Department,  Merchant  Venturers'  Technical  College,  Bristol.     With 

176  diagrams.     Sixth  Edition.     Crown  8vo.     35.  6d. 
BUILDERS'    QUANTITIES,   By  H.   C.  GRUBB,   Lecturer   in  Quantities  to  the 

Beckenham  Technical  Institute.     With  76  diagrams.     Crown  8vo.    43.  6d. 

REPOUSSE  METAL  WORK.    By  A.  C.  HORTH,  Instructor  to  the  London  County 

Council.     With  many  diagrams  and  8  plates.     Crown  8vo.     as.  6d. 
AN  INTRODUCTION  TO  THE  STUDY  OF  TEXTILE  DESIGN.     By  ALFRED  F. 

BARKER,   Head   of   the    Textile   Department,    Bradford  Technical    College. 

With  many  Illustrations  and  design  sheets.     Demy  8vo.     75.  6d. 
HOW  TO  MAKE  A  DRESS.      By  J.    A.    E.    WOOD.     With    45   illustrations. 

Fifth  Edition.     Crown  8vo.     is.  6d. 
MILLINERY,  THEORETICAL  AND  PRACTICAL.    By  CLARE  HILL,  Instructress 

to  the  West  Riding  County  Council  and  Leeds  Education  Committees.     With 

many  Illustrations.     Sixth  Edition.     Crown  8vo.     as. 
INSTRUCTION  IN  COOKERY.     By  A.  P.  THOMPSON,  Instructress  to  the  London 

County  Council.     With  10  Illustrations.     Crown  8vo.     as.  6d. 


A  TEXT-BOOK  OF 
INORGANIC  CHEMISTRY 


BY 


GEORGE  SENTER,  D.SC(LOND.),  PH.D.,  F.I.C. 

•R1NCIPAL    AND    HEAD   OF   THE   CHEMISTRY    DEPARTMENT,    BIRKBECK   COLLEGE,    LONDON 

EXAMINER    IN   CHEMISTRY,    UNIVERSITY   OF    LONDON 

FORMERLY   READER    IN   CHEMISTRY    IN    THE   UNIVERSITY   OF   LONDON 
LECTURER   ON   CHEMISTRY   AT    ST.    MARY'S    HOSPITAL   MEDICAL   SCHOOL 

EXTERNAL   EXAMINER   IN   CHEMISTRY,    UNIVERSITY   OF   BIRMINGHAM 

EXAMINER   IN   CHEMISTRY   TO   THE   ROYAL   COLLEGE   OF   PHYSICIANS   OF    LONDON 

AND   THE    ROYAL   COLLEGE   OF    SURGEONS   OF   ENGLAND 


FIFTH   EDITION 


METHUEN   &   CO.   LTD. 

36    ESSEX    STREET    W.Co 

LONDON 

D.  VAN7  NOSTRA>'D  COMPANY 
YOP.E 


First  Published 
Second  Edition 
Third  Edition  . 
Fourth  Edition 
Fifth  Edition  . 


November  2nd  IQII 
November  1913 
November  fQt6 

October  igiS 

79/9 


PREFACE 

UP  till  about  twenty-five  years  ago,  Inorganic  Chemistry  was 
based  mainly  on  certain  fundamental  laws  and  theories, 
notably  the  Laws  of  the  Conservation  of  Mass  and  of  Energy, 
the  Laws  of  Chemical  Combination,  the  Atomic  Theory,  Avo- 
gadro's  Hypothesis,  and  the  Periodic  System  of  Classifying  the 
Elements.  The  recent  development  of  the  subject  is  char- 
acterized  by  the  discovery  and  applications  of  certain  laws  and 
theories  usually  associated  with  Physical  Chemistry,  more  par- 
ticularly the  principles  of  Chemical  Equilibrium  (including  the 
Law  of  Mass  Action),  the  conception  of  Osmotic  Pressure  and 
its  application  to  the  determination  of  Molecular  Weights  in 
Solution,  and  the  Electrolytic  Dissociation  Theory.  The  newer 
laws  and  theories  are  a  necessary  supplement  to,  and  extension 
of  the  older  principles  established  by  the  labours  of  Boyle, 
Lavoisier,  Richter,  Dalton,  Avogadro  and  Mendeleeff,  but  their 
application  to  the  problems  of  Inorganic  Chemistry  has  to  a 
large  extent  revolutionized  the  study  of  that  subject.  It  is  now 
generally  recognized  that  the  newer  views  have  contributed 
enormously  to  the  development  of  chemistry,  but  in  spite  of  this 
fact  their  general  adoption  into  courses  of  elementary  instruction 
in  this  country  has  been  very  slow.  The  present  book  is  written 
throughout  from  the  modern  standpoint,  and  it  is  hoped  that  it 
may  contribute  in  some  degree  to  the  wider  use  of  the  newer 
principles  at  a  relatively  early  stage  in  chemical  courses. 

Chemistry  is  an  experimental  science,  and  no  adequate  know- 
ledge of  it  can  be  gained  without  practical  work.     The  study  of 

62 


vi     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

a  text-book  should  proceed  simultaneously  with  the  performance 
of  a  number  of  experiments  illustrating  the  general  principles  of 
the  subject.  From  considerations  of  space,  it  has  been  found 
impossible  to  give  full  practical  details  for  experimental  work  in 
this  book,  but  simple  experiments  for  illustrative  purposes  can 
readily  be  devised. 

The  order  of  treatment  of  the  different  parts  of  the  subject 
presented  considerable  difficulty,  and  has  been  decided  by  my 
own  teaching  experience.  One  of  the  guiding  principles  has 
been  that  a  study  of  the  facts  on  which  la^s  and  theories  are 
based  should  precede  the  statement  of  the  laws  and  theories 
themselves.  For  this  reason  the  study  of  oxygen,  hydrogen  and 
water  is  taken  up  at  a  very  early  stage,  and  in  connexion  with 
these  subjects  the  general  properties  of  gases,  liquids  and  solu- 
tions are  discussed,  and  a  brief  account  of  combustion  is  given. 
Chlorine  and  hydrogen  chloride  are  then  considered,  and  in  this 
way  a  sufficient  number  of  facts  are  available  to  illustrate  the 
Laws  of  Chemical  Combination,  the  Atomic  Theory,  the  deduc- 
tion of  Molecular  Formulae  and  the  writing  of  Equations,  which 
are  dealt  with  in  succeeding  chapters.  The  fundamental  distinc- 
tion between  facts  and  theories,  so  often  inadequately  appreciated 
by  the  student,  is  emphasized  by  this  arrangement  of  the  subject- 
matter. 

It  follows  from  the  above  statement  that  the  theoretical  part 
of  the  subject  is  distributed  throughout  the  book.  The  different 
topics  are  dealt  with  as  opportunity  offers,  but  the  treatment  of 
any  particular  subject  is  postponed  till  facts  illustrating  it  have 
been  mentioned.  The  subject  in  question  is  then  considered  as 
fully  as  space  permits — a  mode  of  procedure  which  appears 
preferable  to  dealing  with  it  in  a  fragmentary  manner  in  different 
parts  of  the  book.  The  latter  method  may  usefully  be  employed 
for  lectures,  but  the  former  method  greatly  increases  the  value  of 
the  book  for  purposes  of  reference  and  revision.  As  regards  the 
order,  I  have  also  been  guided  to  some  extent  by  the  historical 
'  development  of  the  science. 


PREFACE  vii 

Subject  to  the  above  considerations,  the  more  important 
branches  of  chemical  theory,  including  the  Principles  of  Chemical 
Equilibrium,  the  determination  of  Molecular  Weights  in  Solu- 
tion, Electrolysis  and  Electrolytic  Dissociation,  are  introduced 
at  a  fairly  early  stage  of  the  work,  and  their  applications 
are  repeatedly  illustrated  in  the  latter  part  of  the  book.  The 
Periodic  System  is  described  after  the  non-metals,  and  serves 
as  the  basis  for  the  discussion  of  the  metals  and  their  com- 
pounds. One  of  the  main  tasks  of  a  teacher  is  so  to  familiarize 
the  student  with  a  general  principle  that  he  learns  to  apply  it 
with  confidence.  The  mere  statement  of  a  principle,  in  the 
expectation  that  the  student  will  straightway  apply  it  in  his  later 
work,  is  useless. 

It  will  be  found  that  the  subjects  usually  included  in  an 
Elementary  Course,  such  as  Stage  I.  of  the  Board  of  Education, 
are  all  dealt  with  in  the  early  part  of  the  book,  and  the 
more  advanced  work  in  the  later  chapters.  Further  division 
of  the  subject  into  stages  leads  inevitably  to  the  result  that 
information  regarding  any  important  element  has  to  be  sought 
for  in  different  parts  of  the  book,  thus  rendering  it  much  less 
useful  for  purposes  of  reference  and  revision. 

The  book  is  designed  for  use  in  University,  Technical  Institute 
and  other  general  classes  on  the  subject,  and  contains  all  that  is 
usually  included  in  a  B.Sc.  Course. 

In  the  preparation  of  the  book,  which  is  based  on  my 
lectures,  a  number  of  larger  works,  more  particularly  Abegg's 
Anorganische  Chemie  and  Ostwald's  Anorganische  Chemie,  have 
been  frequently  consulted;  the  Abstracts  published  by  the 
Chemical  Society  have  also  proved  very  serviceable.  The 
references  to  Physical  Chemistry  in  the  text  are  to  the 
Outlines  of  Physical  Chemistry  (2nd  edition,  1911),  where 
further  information  on  certain  branches  of  the  subject  may 
be  found. 

In  conclusion,  I  desire  to  express  my  sincere  thanks  to  Dr. 
J.  T.  Hewitt,  F.R.S.,  the  General  Editor  of  the  Series,  who 


viii     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

read  the  whole  of  the  manuscript  and  made  many  valuable 
suggestions.  I  must  also  express  my  obligations  to  Prof.  A.  W. 
Porter,  F.R.S.,  and  Dr.  H.  Burrows  for  useful  criticisms,  and 
to  Mr.  T.  J.  Ward,  Mr.  R.  W.  Davies,  and  Dr.  J.  M'Donald 
for  assistance  in  correcting  the  proofs. 

G.  S. 

LUMSDEN,  ABERDEENSHIRE, 
September  1911. 


PREFACE  TO  FOURTH   EDITION 

I^HE  second  edition  of  the  book  differed  from  the  first  in 
containing  a  selection  of  problems  and  questions,  which 
have  been  found  to  contribute  materially  to  its  usefulness.     A 
new  section  on  Acids,  Bases  and  Salts  was  also  added,  and  the 
description  of  technical  processes  was  brought  up  to  date. 

For  the  third  edition,  the  sections  on  Valency  and  on  Radio- 
activity were  rewritten  and  considerably  extended,  and  a  number 
of  new  diagrams,  kindly  prepared  by  my  friend  Mr.  W.  H.  White, 
M.A.,  of  St.  Mary's  Hospital  Medical  School,  have  been 
substituted  with  advantage  for  the  corresponding  figures  in  the 
previous  editions.  For  the  present  edition  the  whole  book  has 
been  carefully  revised,  and  a  number  of  minor  improvements  have 
been  made. 

G.  S. 

HARROW,  September  1917. 


CONTENTS 


CHAP.  PAGE 

I.    INTRODUCTORY  ILLUSTRATIONS         OF          CHEMICAL 

CHANGE            I 

II.    CONSERVATION   OF    MASS  AND   OF   ENERGY — CHEMICAL 

ATTRACTION 8 

III.  THE  CHEMICAL    ELEMENTS             .             .            .             .             •  I? 

IV.  OXYGEN — COMBUSTION 2O 

V.    HYDROGEN — GENERAL    PROPERTIES    OF   GASES     .            .  31 

VI.    WATER — GENERAL    PROPERTIES    OF    LIQUIDS          .             .  51 

VII.    SOLUTION 76 

VIII.    CHLORINE    AND    HYDROCHLORIC    ACID            ...  86 

IX.    LAWS      OF      CHEMICAL      COMBINATION THE      ATOMIC 

THEORY  .  .  .  .  .  .  .  .101 

X.  ^DETERMINATION     OF     ATOMIC      WEIGHTS — COMBINING 
WEIGHTS   AND   CHEMICAL  EQUIVALENTS — FORMULAE 
AND    EQUATIONS — VALENCY    .  .  .  .  .11$ 

XI.    OZONE         AND         HYDROGEN         PEROXIDE THERMO- 
CHEMISTRY                 .             .  133 

XII.    THE    HALOGENS    AND    HALOGEN    ACIDS           .            .            .  149 

XIII.  CHEMICAL    EQUILIBRIUM THERMAL    DISSOCIATION       .  164 

XIV.  OXIDES    AND    OXYGEN    ACIDS    OF    THE    HALOGENS            .  176 
XV.    OSMOTIC      PRESSURE      AND     MOLECULAR      WEIGHT      IN 

SOLUTION        .  .  .  .  .  .  «  .192 


x       A  TEXT-BOOK    OF   INORGANIC    CHEMISTRY 

CHAP.  PAGK 

XVI.   NITROGEN,  THE   ATMOSPHERE   AND   THE   ELEMENTS 

OF    THE    HELIUM    GROUP 2OO 

XVII.   COMPOUNDS    OF    NITROGEN    WITH     HYDROGEN    AND 

WITH    THE    HALOGENS 213 

XVIII.   OXIDES    AND    OXYACIDS    OF    NITROGEN      .           .           .  222 

XIX.   PHOSPHORUS 238 

XX.   ELECTROLYSIS    AND    ELECTROLYTIC    DISSOCIATION  ; 

OXIDES,    ACIDS,    BASES    AND    SALTS         .  ,  .257 

XXI.    SULPHUR,    SELENIUM    AND    TELLURIUM     .           .           .  289 

XXII.   CARBON 324 

XXIII.  COMBUSTION    AND    FLAME 353 

XXIV.  SILICON    AND    BORON              ......  363 

XXV.   CLASSIFICATION  OF  THE  ELEMENTS— THE  PERIODIC 

SYSTEM — GENERAL    PROPERTIES   OF  THE  METALS 

AND    THEIR    COMPOUNDS 378 

XXVI.   THE    ALKALI    METALS 394 

XXVII.    METALS    OF    THE    COPPER    GROUP       .  .  .  .423 

XXVIII.   METALS    OF    THE    ALKALINE    EARTHS          .           .           .  448 

XXIX.    METALS    OF    THE    ZINC    GROUP              ....  460 

XXX.   METALS    OF    THE    ALUMINIUM    GROUP         .           .           .  479 

XXXI.    METALS    OF    THE    TIN    GROUP 492 

XXXII.   METALS    OF    THE    ARSENIC    GROUP     .           .            .           .  •  510 

XXXIII.  METALS    OF    THE    CHROMIUM    GROUP  .  .  -533 

XXXIV.  METALS    OF    THE    MANGANESE    GROUP        .           .           .  543 
XXXV.   METALS    OF    THE    IRON    GROUP             .           .           .  551 

XXXVI.    METALS    OF    THE    PLATINUM    GROUP — VALENCY         .  572 

XXXVII.    RADIO-ACTIVITY 587 

PROBLEMS    AND    QUESTIONS 597 

ANSWERS 606 

INDEX    .                        607 


A    TEXT-BOOK    OF 
INORGANIC    CHEMISTRY 


A  TEXT-BOOK  OF 

INORGANIC    CHEMISTRY 


CHAPTER  I 

INTRODUCTORY— ILLUSTRATIONS  OF 
CHEMICAL   CHANGE 

OF  the  different  branches  of  human  knowledge,  the  study  of 
natural  objects  is  the  most  complex  and  comprehensive  ;  it 
constitutes  the  domain  of  natural  science.  For  the  sake  of  con- 
venience, it  has  been  found  desirable  to  distinguish  between  the 
biological  sciences,  which  are  concerned  with  living  things,  and  the 
physical  sciences,  such  as  astronomy,  geology,  physics  and  chemistry, 
which  are  primarily  concerned  with  the  behaviour  and  properties  of 
non-living  things. 

The  different  branches  of  natural  science  are,  however,  by  no  means 
sharply  marked  off  from  each  other.  The  geologist,  for  instance, 
deals  with  such  questions  as  the  origin  of  a  deposit  of  chalk  and 
its  relative  age  with  regard  to  other  deposits,  but  questions  as  to 
the  composition  of  chalk,  the  nature  of  the  changes  produced  by  heat- 
ing, etc.,  belong  primarily  to  the  provinces  of  physics  and  chemistry. 
In  many  of  the  problems  with  which  the  geologist  has  to  deal,  how- 
ever, a  knowledge  of  the  behaviour  of  his  materials  from  a  chemical 
and  physical  point  of  view  is  indispensable,  so  that  these  branches  of 
knowledge  overlap  in  many  respects.  The  same  is  true  to  a  greater 
or  less  extent  of  all  the  natural  sciences. 

At  this  stage  of  our  work  it  is  not  possible  to  give  a  satisfactory 
definition  of  the  province  of  chemistry,  but  the  following  illustration^ 
will  give  a  preliminary  idea  of  the  nature  of  chemical  changes,  mon 
particularly  with  reference  to  the  broad  distinctions  between  physica 
and  chemical  changes. 
I 


2       A   TEXT-BOOK   OF;  INORGANIC    CHEMISTRY 

Physical  4nd  Chemical  Changes— If  a  stick  of  sulphur  is 
briskiy  rubbed  with  a  ch-y  cloth  it  acquires  the  power  of  attracting  light 
objects,  such  as  small  pieces  of  paper,  and  is  said  to  be  electrified.  If 
olaced  in  hot  water  it  acquires  a  new  property,  that  of  being  able  to 
give  up  heat  to  other  bodies  at  the  ordinary  temperature.  Further, 
if  placed  in  a  test-tube  and  carefully  heated  over  a  Bunsen  flame,  the 
sulphur  melts  to  a  yellowish  or  brownish  liquid  (according  to  the 
degree  of  heating),  but  on  removing  the  tube  from  the  source  of  heat 
the  contents  soon  become  solid  again,  and  then  show  all  the  pro- 
perties of  ordinary  sulphur. 

It  should  be  clearly  realized  that  substances  can  only  be  recognized, 
that  is,  distinguished  from  other  substances,  by  their  properties. 
Thus  the  stick  of  material  which  we. know  as  sulphur  is  characterized 
by  its  colour,  by  its  shape,  by  its  apparent  weight,  the  temperature 
at  which  it  melts,  etc.  Every  substance  has  an  almost  infinite 
number  of  properties,  but  for  purposes  of  recognition  some  pro- 
perties are  much  more  important  and  characteristic  than  others. 
Further  experience  will  show  us  that  it  is  advantageous  to  classify 
the  properties  of  a  substance  as  (i)  characteristic,  those  which 
pertain  to  it  under  all  circumstances,  (2)  non-characteristic,  those 
which  may  alter  without  the  substance  losing  its  identity.  For 
instance,  the  shape  of  a  substance,  such  as  our  sulphur,  is  not  a 
characteristic  property  ;  its  colour,  on  the  other  hand,  is  a  char- 
acteristic property.  We  shall  learn  later  that  the  most  important 
characteristic  property  of  a  substance  is  its  composition. 

In  examining  the  properties  of  a  substance  we  do  not  confine  our 
attention  to  those  which  can  be  observed  directly,  but  extend  it  to 
the  effects  produced  by  altering  the  conditions.  Much  information, 
for  instance,  is  gained  by  noticing  the  effect  of  heat  upon  substances. 

Regarding  the  experiments  with  sulphur  described  above  in  the 
light  of  these  considerations,  we  see  that  only  one  or  two  of  the 
properties — non-characteristic  properties — of  the  sulphur  have  been 
altered,  and  further,  the  changes  are  merely  temporary.  If  left  to 
itself,  the  electrified  sulphur  soon  loses  the  property  of  attracting 
light  objects,  and  the  temperature  of  the  heated  sulphur  soon  falls 
to  that  of  its  surroundings.  Such  changes,  which  are  temporary 
and  affect  only  a  few  of  the  properties  of  a  substance,  are  termed 
physical  changes. 

When  a  little  sulphur  is  placed  on  platinum  foil  and  brought  in 
contact  with  a  flame,  it  catches  fire  and  burns  with  a  bluish  flame, 
giving  rise  to  a  characteristic  sharp  odour ;  in  a  short  time  it  com- 


ILLUSTRATIONS   OF   CHEMICAL   CHANGE         3 

pletely  disappears.  In  this  case  a  much  more  fundamental  change 
has  taken  place,  and  the  change  is  permanent ;  no  substance 
having  any  of  the  original  characteristic  properties  of  sulphur 
remains.  On  the  other  hand  a  new  substance,  characterized  by  a 
sharp,  choking  smell,  is  formed  and  escapes  into  the  atmosphere. 
The  sulphur  has  in  this  case  undergone  a  chemical  change. 

Another  instructive  experiment  is  to  place  successively  in  the  same 
flame  a  piece  of  platinum  wire  and  a  piece  of  magnesium  ribbon. 
The  platinum  wire  becomes  white-hot,  but  on  removal  from  the  flame 
regains  all  its  original  properties,  so  that  the  change  is  a  physical 
one.  The  magnesium  ribbon,  on  the  other  hand,  burns  with  an 
extremely  bright  flame,  and  when  the  change  is  complete  there 
remains  a  new  substance  in  the  form  of  a  white  powder,  which 
differs  in  all  its  properties  from  the  original  metal.  Hence  a 
chemical  change  has  occurred. 

On  the  basis  of  these  experiments,  a  broad  distinction  can  be 
drawn  between  physical  and  chemical  changes.  When  the  change 
is  more  or  less  temporary  and  concerns  only  a  few  of  the  properties 
of  the  substance  it  is  physical;  when,  on  the  other  hand,  new  sub- 
stances, characterized  by  entirely  different  properties,  are  formea, 
the  change  is  a  chemical  one.  It  should  be  carefully  noted  that 
the  chemical  change  of  one  substance  (or  number  of  substances) 
to  others  is  not  gradual,  but  is  perfectly  sharp  and  definite.  This 
is  well  illustrated  by  the  formation  of  the  new  substance  on  burning 
magnesium  ribbon.  This  change  is  usually  incomplete,  especially  it 
performed  in  a  crucible,  and  particles  of  the  white  powder  may  be 
observed  lying  side  by  side  with  particles  of  magnesium.  The  pro- 
perties of  the  two  are  entirely  distinct ;  each  particle  has  either 
changed  completely,  or  not  at  all. 

In  the  extreme  cases  discussed  above,  there  is  no  difficulty  in 
distinguishing  between  physical  and  chemical  changes,  but,  as  we 
shall  learn  later,  the  matter  is  by  no  means  always  so  simple.  It  is 
evident  from  the  foregoing  that  the  essential  point  is  to  be  able  to 
detect  the  formation  of  new  substances,  and,  for  the  sake  of  clear- 
ness, one  or  two  definitions  will  now  be  given. 

Everything  around  us  is  said  to  be  composed  of  matter.  A 
satisfactory  definition  of  such  a  fundamental  corception  as  matter 
cannot  be  given,  but  for  our  present  purpose  matter  may  be  regarded 
as  anything  which  occupies  space  and  has  weight.  A  definite  limited 
portion  of  matter,  such  as  a  piece  of  sulphur,  a  piece  of  granite, 
or  a  knife,  is  termed  a  body  or  thing.  Bodies  differ,  however,  in 


4      A   TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

complexity.  The  material  of  which  the  piece  of  sulphur  is  made  up 
is  homogeneous  to  the  naked  eye,  and  even  under  the  microscope, 
whereas  in  granite  three  constituents  can  readily  be  distinguished, 
a  white  crystalline  part  termed  quartz,  a  gray  portion,  felspar,  and 
nearly  colourless,  lustrous  scales  called  mica.  Each  of  these  three 
components  is  in  itself  uniform  or  homogeneous  to  the  naked  eye 
Such  homogeneous  materials  are  termed  substances.  Whilst  sulphur 
is  made  up  of  a  single  substance,  granite  is  made  up  of  three  sub- 
stances, each  of  which  is  characterized  by  its  special  properties. 

The  important  distinction  between  bodies  and  substances  should 
be  carefully  noted.  Needles,  chains,  hammers,  and  nails  are  different 
bodies,  but  all  may  be  composed  of  the  same  substance,  steel. 

Further  Illustrations  of  Chemical  Change.  Mixtures 
and  Chemical  Compounds— When  a  piece  of  sulphur  is  ground 
in  a  mortar  and  mixed  with  about  its  own  weight  of  iron  filings, 
an  apparently  homogeneous,  grayish  powder  is  obtained.  On  ex- 
amination with  a  microscope,  however,  the  separate  particles  of 
iron  and  sulphur  can  readily  be  detected.  Moreover,  if  a  magnet 
is  held  just  above  the  mixture,  the  iron  can  be  removed,  leaving  the 
sulphur,  and  a  separation  can  also  be  effected  by  shaking  up  the 
mixture  with  a  liquid  called  carbon  disulphide,  which  takes  up 
(dissolves)  the  sulphur,  leaving  the  iron.  It  is  evident  that  we  are 
dealing  with  a  mechanical  mixture  of  iron  and  sulphur,  in  which  both 
substances  retain  all  their  properties  unimpaired. 

If  now  the  mixture  is  placed  in  a  test-tube  and  heated,  the  contents 
of  the  tube  soon  begin  to  glow,  and  this  glowing  increases  and  spreads 
through  the  whole  mass,  even  if  the  tube  is  removed  from  the  flame. 
After  the  tube  has  cooled  to  room  temperature,  a  dark  mass  remains, 
unlike  either  the  iron  or  sulphur.  From  this  mass  the  iron  cannot  be 
abstracted  by  means  of  a  magnet,  nor  can  the  sulphur  be  dissolved  out 
by  means  of  carbon  disulphide.  Even  under  the  highest  power  of  the 
microscope  no  separate  particles  of  iron  and  sulphur  can  be  detected ; 
the  new  substance  is  homogeneous.  Further,  if  a  few  drops  of  dilute 
hydrochloric  acid  are  added  to  a  small  portion  of  the  dark  mass  in  a 
test-tube,  a  gas  with  a  very  disagreeable  odour  is  given  off,  quite 
different  from  that  obtained  by  adding  a  little  of  the  acid  to  the 
mechanical  mixture  of  iron  and  sulphur.  It  is  evident,  therefore, 
that  on  heating  the  mixture  a  chemical  change  has  taken  place  be- 
tween the  iron  and  the  sulphui,  resulting  in  the  formation  of  a  sub- 
stance whose  properties  are  entirely  different  from  those  of  the  original 
substances.  The  new  substance  is  a  definite  chemical  compound 


ILLUSTRATIONS   OF   CHEMICAL   CHANGE        5 

which,  in  allusion  to  its  formation  from  iron  and  sulphur,  is  called 
iron  sulphide. 

The  foregoing  experiments  illustrate  the  more  important  differences 
between  a  mechanical  mixture  and  a  chemical  compound  : 

(1)  A  chemical  compound  has  properties  peculiar  to  itself;  the  pro- 
perties of  a  mixture  are  the  mean  of  the  properties  of  its  constituents. 

(2)  A  chemical  compound  is  homogeneous;  a  mixture  is  usually, 
though  not  invariably,  heterogeneous. 

(3)  A  mixture,  unlike  a  chemical  compound,  can  usually  be  separ- 
ated into  its  constituents  by  mechanical  means. 

To  these  differences  may  be  added  a  fourth,  which  will  be  fully 
discussed  at  a  later  stage  (p.  101). 

(4)  The  composition  of  a  definite  chemical  compound,  unlike  that  of 
a  mixture,  is  constant  and  invariable. 

Another  instructive  chemical  change  will  now  be  described.  A 
small  amount  of  the  red  powder  called  mercuric  oxide  is  cautiously 
heated  in  a  test-tube  until  it  turns  black.  If  at  this  point  the  tube 
is  removed  from  the  flame,  the  powder  regains  its  original  colour  and 
other  properties  on  cooling,  and  the  change  in  question  is,  therefore,  a 
physical  one.  If,  however,  the  heating  is  continued,  it  will  be  noticed 
that  a  mirror  soon  begins  to  form  on  the  cooler  upper  part  of  the 
tube,  and,  further,  if  a  glowing  splinter  is  inserted  in  the  tube  it  will 
burst  into  flame.  The  tube  is  now  removed  from  the  source  of  heat, 
and  the  mirror  rubbed  with  a  glass  rod,  when  small  shining  globules 
will  be  formed,  characteristic  of  mercury.  The  property  of  causing  a 
glowing  splinter  to  burst  into  flame  indicates  the  presence  of  a  colour- 
less gas,  called  oxygen.  In  this  case,  therefore,  under  the  influence 
of  heat,  the  red  powder  has  given  rise  to  two  new  substances,  the  liquid 
metal  mercury  and  the  colourless  gas  oxygen. 

An  experiment  of  a  different  type  will  now  be  described.  When  a 
few  crystals  of  ordinary  sodium  carbonate  (washing  soda)  are  shaken 
up  with  water,  they  soon  disappear  and  form  a  homogeneous  mixture 
with  the  water.  Such  a  homogeneous  mixture  is  called  a  solution, 
and  the  sodium  carbonate  is  said  to  have  dissolved  in  the  water.  A 
solution  in  water  of  the  substance  called  calcium  chloride  is  prepared 
in  the  same  way.  On  mixing  these  clear  solutions  a  white  substance 
is  produced,  which,  in  course  of  time,  partially  settles  to  the  bottom  of 
the  mixture.  The  solid  substance  may  be  separated  from  the  remainder 
of  the  mixture  by  pouring  the  contents  of  the  beaker  on  a  filter-paper 
supported  in  a  glass  funnel,  as  shown  in  Fig.  I.  Filter-paper  is  a  porous 
form  of  paper  which  readily  allows  liquids  to  pass  through,  but  retains 
solid  substances.  The  solid  residue  on  the  filter-paper  may  be  shown 


6      A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

to  have  all  the  properties  of  ordinary  chalk  (calcium  carbonate).  It 
is  still  possible,  however,  that  another  substance  may  have  been 
formed  by  interaction  of  sodium  carbonate  and  calcium  chloride 
in  solution,  but  has  remained  dissolved  in  the  water.  The  simplest 
way  of  testing  this  possibility  is  to  heat  the  solution  which  has  passed 
through  the  filter-paper — known  as  the  filtrate — until  all  the  water  is 
driven  off.  Finally  a  residue  is  left  which  is  salt  to  the  taste,  and  has 
all  the  properties  of  common  salt  (sodium  chloride).  The  change  in 


FIG.  i. 

this  case  is,  therefore,  a  somewhat  complicated  one,  two  substances 
entering  into  chemical  combination  and  giving  rise  to  two  new  sub- 
stances. 

The  foregoing  experiments,  which  should  be  performed  by  the 
student,  illustrate  the  three  most  important  types  of  chemical  change, 
which  are  as  follow  : — 

(i)  Chemical  Combination,  when  two  or  more  substances  unite  to 
form  a  single  substance.  Example,  iron  and  sulphur.  If  the  symbols 
A  and  B  represent  the  two  substances,  this  type  of  chemical  change 
may  be  represented  thus — 

A  +  B->AB, 


ILLUSTRATIONS   OF   CHEMICAL   CHANGE         7 

where  the  approximation  of  the  letters  on  the  right  of  the  arrow 
indicate  that  the  substances  are  chemically  combined. 

(2)  Chemical  Decomposition,  when  a  chemical  compound  splits  up 
into  two  or  more  substances.  The  decomposition  of  mercuric  oxide 
by  heat  is  a  good  illustration  of  this  type  of  change.  In  symbols 
it  is  represented  thus  — 


(3)  Double  Decomposition^  when  two  (or  more)  substances  interact 
to  form  two  (or  more)  new  substances.  Example,  sodium  carbonate 
and  calcium  chloride,  as  just  described.  A  simple  case  of  double 
decomposition  is  represented  symbolically  thus  — 

A  +  B-^C  +  D, 

where  A  and  B  stand  for  the  reacting  substances  and  C  and  D  for 
the  products. 

One  important  point  in  connexion  with  the  chemical  changes 
described  is  that  the  reacting  substances  must  be  in  actual  contact 
before  any  reaction  occurs.  This  is  an  important  characteristic  of  all 
chemical  changes. 

Elements  and  Chemical  Compounds  —  The  foregoing  ex- 
amples show  that  under  certain  conditions  a  chemical  compound  can 
be  split  up  into  simpler  substances,  as  in  the  case  of  mercuric  oxide. 
It  is  natural  to  inquire  whether,  by  further  heating  or  otherwise,  these 
substances  can  be  split  up  into  anything  still  simpler.  So  far,  this  has 
not  been  done  ;  both  mercury  and  oxygen  have  up  to  the  present 
resisted  all  attempts  at  further  simplification.  Substances  of  this  type 
are  termed  elements.  The  accepted  view  of  the  constitution  of  the  uni 
verse  is  that  it  is  made  up  of  a  large  number  of  different  kinds  of  matter, 
the  elements,  and  no  method  is  known  by  which  any  element  can  be 
further  simplified  or  can  be  converted  into  another  element  at  will. 
Chemical  compounds  are  made  up  of  two  or  more  elements  in  chemical 
combination.  A  binary  compound  is  one  which  contains  two  elements 
only.  Many  of  the  common  metals,  such  as  iron,  lead,  tin,  and  zinc 
are  elements,  as  are  sulphur,  phosphorus,  and  charcoal  (carbon). 

Up  to  the  present  about  eighty  elements  have  been  discovered. 
They  are  enumerated,  and  some  of  their  more  important  properties 
briefly  considered,  in  Chapter  III. 

1  As  will  be  pointed  out  in  detail  later  (chap,  xxxvii.)  it  has  been  found  quite 
recently  that  certain  elements,  more  particularlv  radium,  have  the  power  of  de- 
composing spontaneously  with  ultimate  production  of  other  elements,  but  up  to  the 
present  no  method  of  initiating  or  controlling  such  changes  has  been  discovered. 


CHAPTER   II 

CONSERVATION   OF    MASS   AND  OF   ENERGY- 
CHEMICAL  ATTRACTION 

Conservation  of  Weight  in  Chemical  Changes— So  far, 
^-^  we  have  considered  chemical  changes  from  the  qualitative 
point  of  view  only,  but  it  is  also  necessary  to  consider  them  from 
the  point  of  view  of  the  relative  amounts  of  the  reacting  substances. 
As  a  preliminary  to  this  inquiry,  it  is  first  necessary  to  ascertain 
whether  chemical  changes  are  associated  with  changes  in  weight  of 
the  reacting  substances.  The  indispensable  instrument  in  such  in- 
vestigations is  the  balance,  by  means  of  which  the  mass,  or  quantity 
of  matter,  is  determined. 

The  balance  (Fig.  2)  consists  essentially  of  a  long  beam,  supported 
on  a  knife-edge  at  its  middle  point,  and  provided  with  scale-pans 
supported  on  knife-edges  at  either  extremity  of  the  beam.  The 
substance  of  unknown  weight  is  put  on  one  scale-pan  and  known 
weights  on  the  other  until  the  beam  is  exactly  horizontal,  as  indicated 
by  a  pointer  in  front  of  a  scale. 

The  balance  does  not  indicate  directly  the  mass  or  quantity  of  matter 
in  a  body,  but  its  weight,  that  is,  the  force  with  which  it  is  attracted 
towards  the  earth.  The  weight  of  a  substance  is  the  product  of  its 
mass  and  the  force  of  gravity  at  the  place  of  observation.  Hence,  as 
at  any  one  place  the  weight  and  the  mass  bear  a  constant  ratio  to 
one  another,1  two  bodies  of  the  same  weight  have  equal  masses  or 
contain  equal  quantities  of  matter. 

We  are  now  in  a  position  to  investigate  the  question  as  to  whether 
there  is  any  alteration  in  mass  (or  in  weight)  when  substances  enter 
into  chemical  combination.  Some  well-known  experiments  appear  at 
first  sight  to  show  that  such  changes  in  weight  actually  occur. 

Into  an  ordinary  crucible  provided  with  a  lid  some  pieces  of 
magnesium  ribbon  are  placed,  and  the  crucible  and  contents  weighed. 
The  covered  crucible  is  then  strongly  heated  over  a  Bunsen  flame, 

1  The  ratio  between  weight  and  mass,  in  other  words  the  force  of  gravity,  is 
different  at  different  parts  of  the  earth's  surface. 


FIG.  2. 


CONSERVATION  OF  MASS  AND  OF  ENERGY  9 

the  lid  being  raised  occasionally  so  as  to  admit  air.  When  the 
glowing  of  the  metal  has  practically  ceased,  the  crucible  and  contents 
are  allowed  to  cool  and  again 
weighed.  If  the  experiment 
has  been  carefully  performed 
it  will  be  found  that  there  is 
a  gain  in  weight.  This,  how- 
ever, does  not  show  conclu- 
sively that  the  products  of 
the  chemical  change  weigh 
more  than  the  substances 
which  entered  into  reaction, 
as  in  the  course  of  the  heat- 
ing something  may  have 
been  taken  up  from  the  air. 
When  a  candle  burns  in 
the  air  it  slowly  disappears, 
and  in  this  case  it  seems  as  if  the  chemical  change  is  attended  by  a  loss 
of  weight.  There  is,  however,  the  possibility  that  something  may  in 
this  case  be  passing  into  the  air,  thus  escaping  being  weighed,  and 
this  may  be  shown  to  be  the  case  by  the  following  experiment.  A 
piece  of  candle  is  fixed  on  a  piece  of  cork  cut  so  as  to  fit  the  bottom 
of  a  glass  cylinder  (lamp  glass),  the  cork  being  provided  with  a  large 
number  of  holes  for  the  admission  of  air.  In 
the  upper  part  of  the  cylinder  a  piece  of  wire 
gauze  is  supported,  and  the  space  above  it  is 
filled  with  sticks  of  a  substance  known  as  caustic 
soda,  which,  as  we  shall  see  later,  has  the  property  of 
taking  up  the  substances  formed  when  a  candle  burns  in 
the  air  (Fig.'  3).  The  whole  arrangement  is  then  placed 
on  one  pan  of  a  balance  and  weights  just  sufficient  to 
bring  the  pointer  to  the  middle  of  the  scale  placed  on 
the  other.  As  the  candle  burns,  it  will  be  observed  that 
that  side  of  the  balance  is  slowly  depressed,  showing  a 
gain  in  weight.  The  same  remark  applies  here,  how- 
ever, as  in  the  magnesium  experiment ;  the  increase  in 
weight  may  be  accounted  for  by  the  absorption  of  some- 
thing from  the  air. 

It  is  evident  that  in  order  to  ensure  that  all  the  reacting 
substances  and  all  the  products  are  taken  into  account,  the  reaction 
must  be  carried  out  in  a  closed  space.  A  convenient  chemical  change 


FIG.  3. 


lo     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

for  our  purpose  is  the  action  of  sodium  carbonate  on  calcium  chloride 
in  aqueous  solution,  which  has  already  been  described  in  another  con- 
nexion (p.  5).     A  wide-mouthed  glass  bottle   (Fig.  4)   contains  a 
small  amount  of  an  aqueous  solution  of  sodium  carbon- 

Qate,  and  a  small  tube,  containing  an  aqueous  solution  of 
calcium  chloride,  is  also  placed  in  the  bottle  in  such  a 

| •{       way  that  no  mixing  of  the  solutions  takes  place ;  the 

^     bottle  is  then  well  corked  and  the  whole  weighed.    The 
bottle  is  then  inclined  in  such  a  way  that  the  solutions 
mix  and    double    decomposition    (p.    7)  takes   place. 
On  weighing  again,  it  will  be  found  that  the  weight  is 
unaltered.     Experiments  with  other  substances  lead  to 
the  same  conclusion ;  provided  that  all  the  reacting  sub- 
stances and  all  the  products  of  reaction  are  taken  into 
FIG.  4.         account,  there  is  no  change  in  weight  as  the  result  of  a 
chemical  change.     The  short  statement  of  these  results 
is  called  the  LAW  OF  THE  CONSERVATION  OF  MASS,  and  should  be 
remembered  in  the  following  form  : 

When  a  chemical  change  occurs^  the  total  weight  (or  mass)  of  the 
reacting  substances  is  equal  to  the  total  weight  (or  mass)  of  the 
products.  The  law  holds  for  physical  as  well  as  for  chemical 
changes. 

The  above  statement  is  sometimes  called  the  law  of  the  conservation 
of  matter,  and  is  regarded  as  indicating  that  the  quantity  of  matter  in 
the  universe  cannot  be  altered  in  consequence  of  chemical  (or  any 
other)  changes  ;  in  other  words,  matter  is  indestructible.  The  defini- 
tion in  italics  is,  however,  to  be  preferred,  as  it  is  purely  experi- 
mental, whereas  the  latter  statement  introduces  difficulties  with 
reference  to  the  proper  definition  of  matter. 

The  law  of  the  conservation  of  weight,  which  was  firmly  established 
by  the  investigations  of  Lavoisier  towards  the  end  of  the  eighteenth 
century,  is  of  the  most  fundamental  importance  for  chemistry.  Al- 
though it  may  seem  at  first  sight  that  the  law  is  self-evident,  it  must 
be  clearly  realized  that  the  law  is  a  purely  experimental  one,  and  can 
only  be  regarded  as  established  within  the  limits  of  the  unavoidable 
errors  of  the  experiments  which  have  been  made  to  test  it. 

Landolt-s  Experiments  on  Conservation  of  Weight — 
The  most  comprehensive  series  of  experiments  undertaken  with 
the  object  of  testing  the  validity  of  the  law  in  question  are  due  to 
Landolt.  In  the  majority  of  Landolt's  experiments  glass  tubes  with 
two  limbs  were  used.  The  reacting  substances  were  placed  separately 


CONSERVATION   OF   MASS   AND   OF   ENERGY      n 

one  in  each  limb,  the  tube  carefully  sealed  and  weighed.  The  tube 
was  then  inclined  so  as  to  mix  the  two  solutions,  and  when  the 
chemical  action  was  complete  and  the  temperature  had  fallen  to  that 
of  the  atmosphere,  was  again  weighed.  A  special  form  of  balance 
was  used,  capable  of  detecting  extremely  small  differences  of  weight. 
The  final  result  of  Landolt's  experiments,  which  extended  over  more 
than  twenty  years,  is  that  the  law  of  the  conservation  of  weight  has 
been  verified  to  a  very  high  degree  of  approximation ;  in  no  case 
were  the  very  slight  deviations  observed  greater  than  the  possible 
experimental  error. 

The  Atmosphere  and  Burning— It  now  remains  to  reconcile 
the  fact  that  an  increase  of  weight  is  observed  when  magnesium  and 
a  candle  burn  in  the  air  with  the  law  of  the  conservation  of  weight. 
The  observed  results  could  be  satisfactorily  accounted  for  if  some- 
thing is  taken  up  from  the  atmosphere  during  the  process  of  burning, 
so  that  the  respective  increases  in  weight  observed  with 
the  magnesium  and  with  the  candle  are  exactly  balanced 
by  the  loss  of  weight  suffered  by  the  atmosphere.     The 
detailed  proof  of  the  validity  of  this  suggestion  can  only 
be  given  at  a  later  stage  (p.  28),  but  it  can  readily  be 
shown  that  the  air  has  weight.    A  large  glass  globe 
(Fig.  5)  provided  with  a  stopcock  (a  brass  stopcock  is 
suitable)  is  placed  on  one  pan  of  the  balance  and  weights 
added  to  the  other  pan  until  it  is  in  equilibrium.     Part        FJG>  ^. 
of  the  air  is  then  pumped  out,  the  stopcock  is  closed, 
and  after  replacing  on  the  pan  the  globe  and  contents  prove  to  be 
considerably  lighter  than  before. 

As  will  be  shown  in  detail  later,  the  constituent  which  is  taken  up 
from  the  atmosphere  during  burning  is  oxygen,  a  gas  obtained,  as 
already  mentioned,  by  the  action  of  heat  on  mercuric  oxide.  There 
is  now  no  difficulty  in  understanding  the  burning  of  magnesium  and 
of  a  candle  in  the  atmosphere.  In  the  former  case  the  magnesium 
enters  into  chemical  combination  with  the  oxygen  of  the  atmosphere, 
forming  a  white  powder  which  necessarily  weighs  more  than  the 
magnesium  first  taken.  Similarly,  the  constituents  of  the  candle 
combine  with  oxygen  during  burning,  giving  rise  to  gaseous  products 
which  are  taken  up  by  the  sticks  of  caustic  soda.  The  arrangement 
increases  in  weight  during  the  burning,  to  an  extent  determined  by 
the  weight  of  the  oxygen  absorbed. 

The  Conservation  of  Energy.  Chemical  Energy— It 
was  shown  in  describing  the  influence  of  heat  on  a  mixture  of  iron 


12     A   TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

and  sulphur  (p.  4)  that  when  chemical  combination  has  properly 
started  the  reaction  proceeds  of  itself,  and  is  accompanied  by  a 
considerable  evolution  of  heat  and  light.  Further  investigation  has 
shown  that  chemical  changes  are  invariably  accompanied  by  heat 
changes ;  in  some  cases  heat  is  given  out,  in  other  cases  it  is 
absorbed. 

Heat  is  a  form  of  energy.  Energy  may  be  defined  as  that  property 
of  a  body  which  diminishes  when  work  is  done  by  the  body,  and  its 
diminution  is  measured  by  the  amount  of  work  done  by  the  body. 
Other  forms  of  energy,  besides  heat,  are  potential  energy,  kinetic 
energy,  electrical  energy,  and  radiant  energy.  For  a  full  discussion 
of  this  subject  a  text-book  on  physics  should  be  consulted. 

One  kind  of  energy  can  be  transformed  into  another.  Thus  if  the 
energy  of  a  falling  weight  be  used  to  drive  a  stirrer  in  water,  the 
water  becomes  hot,  and  the  potential  energy  which  the  weight  possessed 
before  it  began  to  fall  has  been  converted  into  heat.  Further,  if  a 
Bunsen  flame  is  placed  under  a  hot-air  engine,  the  latter  is  set  in 
motion,  so  that  the  heat  in  this  instance  is  partially  transformed  to 
mechanical  (kinetic)  energy.  The  reader  will  be  able  to  supply 
many  other  illustrations  of  transformation  of  energy  from  his  own 
experience. 

Now  it  has  been  found  that  when  a  certain  amount  of  one  form  of 
energy  disappears,  an  equivalent  amount  of  another  form  of  energy 
makes  its  appearance ;  in  other  words,  energy  can  neither  be  created 
nor  destroyed.  This  result,  which  is  purely  experimental,  is  termed 
the  LAW  OF  THE  CONSERVATION  OF  ENERGY,  and  may  be  expressed 
as  follows  :  The  energy  of  an  isolated  system  is  constant^  that  is,  it 
cannot  be  altered  in  amount  by  interactions  between  the  parts  of  the 
system.  By  an  isolated  system  we  mean  one  which  is  neither  receiv- 
ing energy  from  outside  nor  giving  up  energy  to  its  surroundings. 
If,  for  example,  two  substances  capable  of  entering  into  chemical 
combination  are  contained  in  a  closed  vessel,  through  the  walls  of 
which  no  energy  enters  or  passes  out,  the  total  energy  inside  is  the 
same  before  and  after  chemical  combination. 

We  may  assume  that  the  universe  is  made  made  up  of  two  things, 
and  of  two  things  only,  matter  and  energy,  neither  of  which  can  be 
altered  as  regards  total  quantity,  although  they  may  be  altered 
in  form. 

During  the  combination  of  iron  and  sulphur  a  considerable  amount 
of  heat  is  given  out,  as  already  mentioned,  and  therefore,  according  to 
the  law  of  the  conservation  of  energy,  an  equivalent  amount  of  another 


CONSERVATION   OF   MASS   AND   OF   ENERGY      13 


form  of  energy  must  have  disappeared.  The  latter  form  is  con- 
veniently called  chemical  energy,  and  we  state  that  iron  and  sulphur, 
in  the  uncombined  condition,  possess  a  considerable  store  of  chemical 
energy,  part  of  which  is  transformed  into  heat  and  light  when  they 
enter  into  chemical  combination.  It  must  not  be  assumed  that  iron 
sulphide  has  no  chemical  energy— we  know,  as  a  matter  of  fact,  that 
it  has — but  it  possesses  less  energy  than  the  free  elements  before 
combination. 

Chemical  energy  may,  however,  readily  be  transformed  into  other 
kinds  of  energy  than  heat.  This  may  be  shown  very  readily  by 
means  of  the  action  of  a  mixture  of  sulphuric  acid  and  water  on 
metallic  zinc.  When  the  zinc  and  dilute  acid  are  brought  together  in 
a  test-tube  a  very  vigor- 
ous evolution  of  gas  takes 
place,  and  the  tempera- 
ture of  the  solution  in- 
creases considerably,  as 
shown  by  a  thermometer 
placed  in  the  tube.  The 
same  change  may  now  be 
brought  about  in  another 
way.  A  plate  of  zinc  and 
a  plate  of  copper  are 
joined  by  means  of  wires 
to  an  instrument  (Fig.  6) 
known  as  a  galvanometer, 
the  needle  of  which  moves  along  the  scale  when  an  electric  current 
passes  through.  The  metals,  without  being  allowed  to  touch,  are 
dipped  into  a  mixture  of  sulphuric  acid  and  water,  when  it  will  be 
found  that  the  zinc  dissolves  and  simultaneously  an  electric  current 
passes  through  the  galvanometer.  In  this  case  the  chemical  energy 
which  disappears  when  zinc  and  sulphuric  acid  enter  into  chemical 
combination  appears,  in  part  at  least,  as  electrical  energy. 

The  conversion  of  chemical  into  other  forms  of  energy  is  of  the 
utmost  commercial  importance.  Our  most  important  source  of 
energy  is  the  burning  of  coal,  a  process  in  which  the  materials  of 
the  coal  combine  with  the  oxygen  of  the  atmosphere,  the  enormous 
store  of  chemical  energy  which  the%  substances  contain  before  com- 
bination thus  becoming  available  (p.  352). 

Conversion  of  Heat  and  of  Electrical  Energy  into 
Chemical  Energy— The  converse  of  the  transformations  of 


FIG.  6. 


i4     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


energy  described  above  can  also  be  readily  performed.  Many 
compounds  can  be  split  up  by  heat  into  simpler  substances  which 
possess  more  chemical  energy  than  the  original  compounds  ;  in  such 
cases  heat  has  been  partially  transformed  into  chemical  energy. 
The  action  of  heat  on  mercuric  oxide  is  a  process  of  this  nature,  as 
free  mercury  and  free  oxygen  possess  more  energy  than  the  corre- 
sponding amount  of  mercuric  oxide. 

In  a  similar  way,  an  electric  current  may  be  used  to  split  up 
complex  substances  into  simpler  ones  possessing  more  chemical 
energy  ;  in  such  a  case,  electrical  energy  has  been  transformed  into 
chemical  energy.  On  account  of  the  great  im- 
portance of  this  principle,  we  will  illustrate  it  by 
an  account  of  the  decomposition  of  water  by  the 
agency  of  an  electric  current.  Up  till  about  1  30 
years  ago  water  was  regarded  as  a  chemical 
element,  and  the  experiment  now  to  be  described 
is  one  of  the  most  convenient  for  the  demonstra- 
tion of  its  compound  nature. 

The  apparatus  used  for  the  purpose,  known  as 
a  voltameter,  is  represented  in  Fig.  7.  It  con- 
sists essentially  of  a  tube  bent  in  the  form  of  the 
letter  U  (so-called  U-tube)  joined  at  the  lower 
part  to  a  longer  tube  entling  in  a  bulb  B.  In  the 
lower  part  of  the  two  limbs  of  the  U-tube  are  two 
platinum  plates,  a,  a,  connected  with  platinum 
wires  which  are  sealed  through  the  glass  and 
can  be  joined  (usually  by  copper  wires)  to  the 
positive  and  negative  poles  respectively  of  a 
battery.  In  order  to  perform  an  experiment, 
the  stopcocks  at  the  upper  ends  of  the  two  limbs 
are  opened  and  dilute  sulphuric  acid  l  is  poured  into  B  till  it  rises  to 
the  level  of  the  stopcocks,  which  are  then  closed.  As  soon  as  con 
nexion  is  made  with  the  battery,  bubbles  of  gas  are  given  off  from 
each  platinum  plate,  and  rise  to  the  upper  part  of  the  two  tubes. 
After  the  current  has  passed  for  some  time,  it  will  be  observed  that 
the  volume  of  gas  which  has  collected  above  the  platinum  plate 
connected  with  the  negative  pole  of  the  battery  is  about  double  that 
in  the  other  limb.  0 

The  gases  are  colourless,  and  can  easily  be  shown  to  be  both 
odourless  and  tasteless.     The  gas  present  in  larger  proportion  can 
i  Water  containing  a  little  sulphuric  acid. 


FIG.  7. 


CONSERVATION   OF   MASS   AND   OF   ENERGY      15 

be  collected  in  a  small  test-tube  by  inverting  the  latter  over  the  top 
of  the  voltameter  tube  and  cautiously  opening  the  tap.  When  a 
light  is  put  to  the  mouth  of  the  tube  the  gas  takes  fire  and  burns 
with  an  almost  colourless  flame.  By  means  of  these  and  other 
tests,  it  can  be  recognized  as  hydrogen,  a  gas  which  can  be  made 
in  quantity  by  more  convenient  methods  (p.  31).  It  can  further  be 
shown,  from  its  property  of  igniting  a  glowing  splinter  and  by  other 
tests,  that  the  gas  in  the  other  tube  is  oxygen,  the  preparation  of 
which  from  mercuric  oxide  has  already  been  described  (p.  5). 

The  effect  of  passing  an  electric  current  through  acidulated  water 
is  therefore  very  remarkable,  inasmuch  as  at  one  plate — the  negative 
plate — only  hydrogen  is  given  off,  at  the  other  plate  only  oxygen 
is  given  off,  and  the  volume  of  the  hydrogen  is  about  double  that 
of  the  oxygen.  We  shall  see  later  on  that  a  mixture  of  hydrogen 
and  oxygen  can  be  caused  to  combine  with  formation  of  water,  and 
in  this  process  a  large  amount  of  heat  is  given  out  (p.  37).  The 
process  just  described  represents  the  reverse  of  this,  inasmuch  as 
water  has  been  split  up  into  two  component  elements,  hydrogen  and 
oxygen,  and  a  large  amount  of  electrical  energy  has  been  absorbed  in 
the  process.  There  is  evidence,  which  cannot,  however,  be  given 
here,  that  it  is  really  the  water  which  is  split  up,1  but  a  little  sulphuric 
acid  must  also  be  added,  as  pure  water  does  not  conduct  the  electric 
current.  Alkali  may  be  used  in  place  of  acid. 

The  platinum  plate  connected  with  the  positive  pole  of  the  battery 
is  termed  the  positive  pole,  positive  electrode  or  anode ;  the  plate 
connected  with  the  negative  pole  of  the  battery  is  called  the  nega- 
tive pole,  negative  electrode  or  cathode.  The  process  of  splitting 
up  a  chemical  compound  by  means  of  the  electric  current  is  called 
electrolysis  (that  is,  splitting  up  by  means  of  electricity).  The  nature 
of  the  chemical  changes  occurring  in  the  electrolysis  of  water  will 
be  dealt  with  later,  under  the  heading  Electrolysis  (p.  263). 

The  Cause  of  Chemical  Change— It  has  already  been  pointed 
out  that  when  platinum  wire  is  heated  in  the  air  it  suffers  no  chemical 
change,  whilst  magnesium  ribbon  under  the  same  circumstances 
gives  rise  to  a  new  compound,  and  it  is  natural  to  inquire  into  the 
reasons  for  this  difference  of  behaviour.  Chemists  are  accustomed 
to  state  that  there  is  a  certain  attraction — the  so-called  chemical 
affinity — between  magnesium  and  one  of  the  constituents  of  the 
atmosphere  (the  oxygen)  which  comes  into  play  when  the  magnesium 
is  heated,  and  leads  to  chemical  combination,  whereas  there  is  little 
or  no  chemical  affinity  or  chemical  attraction  between  platinum  and 
i  Compare  Physical  Chemistry,  4th  Edition,  p.  400. 


16     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

oxygen  (p.  574).  Similarly,  the  mercury  and  oxygen  in  mercuric 
oxide  are  held  together  by  chemical  attraction,  and  the  substance 
has  to  he  raised  to  a  high  temperature  in  order  to  overcome  this 
attraction  and  liberate  the  elements. 

Very  little  is  known  with  certainty  as  to  the  nature  of  this  so-called 
affinity  or  attraction.  The  term  "  affinity "  denotes  likeness,  and 
was  introduced  into  chemistry  at  a  time  when  it  was  thought  that 
chemical  combination  took  place  most  readily  between  substances 
of  like  properties.  It  will  be  shown  later,  however,  that  this  view  is 
erroneous  ;  the  most  stable  compounds  are  formed  by  combination 
of  substances  of  unlike  properties. 

Summary.  Characteristics  of  Chemical  Change— The 
first  two  chapters  have  been  mainly  devoted  to  a  consideration  of  the 
nature  of  chemical  change.  It  will  be  useful  to  summarize  here  the 
more  important  characteristics  of  chemical  change. 

(1)  As  the  result  of  a  chemical  change  new  substances  make  their 
appearance.     These  new  substances  are  recognized  by  their  pro- 
perties, more  particularly  by  their  composition.     The  composition  of 
a  substance,  that  is,  the  elements  of  which  it  is  built  up  and  the 
proportion  in  which  they  are   present,   is  its  most   characteristic 
property. 

(2)  When  a  chemical  change  takes  place,  the  total  weight  of  the 
products  is  equal  to  the  total  weight  of  the  reacting  substances  (Law 
of  Conservation  of  Mass). 

(3)  A  chemical  change  is  always  attended  by  the  evolution  or 
absorption  of  heat,  or  more  generally,  the  total  chemical  energy  of 
the  final  products  always  differs  from  that  of  the  reacting  substances. 

(4)  Chemical  changes  occur  only  between  substances  which  are 
in  actual  contact. 

(5)  A  fifth  characteristic  will  be  mentioned  here  for  the  sake  of 
completeness,  and  will  be  fully  considered  in  later  chapters.     It  is 
that  chemical  changes  always   occur  between  definite  weights  (or 
volumes,  in  the  case  of  gases)  of  the  reacting  substances.     From  this 
it  follows  that  the  composition  of  a  definite  chemical  compound  is 
constant,  no  matter  how  it  is  prepared,  or  what  is  its  source.     These 
characteristics  of  chemical   change  are  important,  and  should  be 
carefully  remembered. 


CHAPTER  III 
THE  CHEMICAL  ELEMENTS 

WE  have  learnt  in  the  previous  chapters  that  many  of  the  sub- 
stances with  which  we  are  acquainted  can  be  split  up  into 
simpler  substances  by  various  methods,  for  example  by  the  action  of 
heat  or  of  electrical  energy.  Thus  from  a  definite  quantity  of 
mercuric  oxide  by  the  action  of  heat,  two  substances,  mercury  and 
oxygen,  are  obtained,  each  of  which  weighs  less  than  the  original 
oxide,  whilst  the  sum  of  their  weights  is  equal  to  that  of  the  original 
oxide.  So  far,  no  method  has  been  discovered  by  means  of  which 
mercury  and  oxygen  can  be  split  up  into  "simpler"  substances;  in 
other  words,  it  has  not  been  found  possible  to  separate  a  definite 
amount  of  either  of  those  substances  into  two  or  more  other  sub- 
stances, each  weighing  less  than  the  original  substance.  For  this 
reason  mercury  and  oxygen  are  classed  as  elements.  It  must  be  care- 
fully noted  that  elements  are  substances  which  so  far  have  resisted 
all  attempts  at  decomposition,  but  it  does  not  in  the  least  follow  that 
they  are  really  undecomposable  ;  it  is,  in  fact,  highly  probable  that  in 
the  near  future  methods  of  simplifying  at  least  some  of  the  elements 
will  be  discovered. 

The  above  definition  of  an  element  is  ciue  to  Boyle  (1661),  and 
was  later  adopted  by  Lavoisier.  It  has  proved  extremely  serviceable, 
as  many  compounds  which  in  Lavoisier's  day  were  classified  as 
elements  have  since  been  shown  to  be  chemical  compounds.  Bearing 
this  in  mind,  it  might  be  supposed  that  some  substances  accepted  as 
elements  at  the  present  day  are  comparatively  simple  compounds  of 
the  same  type  as  water,  which  could  be  split  up  by  a  more  energetic 
use  of  the  means  now  at  our  disposal.  We  shall  find  later,  however, 
that  this  suggestion  is  highly  improbable.  While  it  may  be  accepted 
that  in  one  sense  the  elements  are  complex,  their  complexity  is  of  a 
different  order  from  that  of  ordinary  chemical  compounds  (p.  382). 

At  present  about  eighty  elements  are  known.  The  exact  number 
cannot  be  definitely  stated,  as  some  of  them  are  present  only  in  very 
small  proportion  on  the  earth,  and  considerable  difficulties  are  met 
2  '7 


1 8     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

with  in  obtaining  them  in  a  pure  condition.  The  names  of  the 
elements  now  recognized  by  the  International  Committee  on  Atomic 
Weights  are  given  in  alphabetical  order  in  the  following  table.  The 
more  important  elements  are  printed  in  ordinary  type,  the  less  im- 
portant in  italics. 


Aluminium 

Fluorine 

Molybdenum 

Silver 

Antimony 

Gadolinium 

Neodymium 

Sodium 

Argon 

Gallium 

Neon 

Strontium 

Arsenic 

Germanium 

Nickel 

Sulphur 

Barium 

Glucinum  (beryllium) 

Nitrogen 

Tantalum 

Bismuth 

Gold 

Osmium 

Tellurium 

Boron 

Helium 

Oxygen 

Terbium 

Bromine 

Hydrogen 

Palladium 

Thallium 

Cadmium 

Indium 

Phosphorus 

Thorium 

Caesium 

Iodine 

Platinum 

Thulium 

Calcium 

Iridium 

Potassium 

Tin 

Carbon 

Iron 

Praseodymium 

Titanium 

Cerium 

Krypton 

Radium 

Tungsten 

Chlorine 

Lanthanum 

Rhodium 

Uranium 

Chromium 

Lead 

Rubidium 

Vanadium 

Cobalt 

Lithium 

Ruthenium 

Xenon 

Columbium 

Lutecium 

Samarium 

Ytterbium  (Neoytterbium) 

Copper 

Magnesium 

Scandium 

Yttrium 

Dysprosium 

Manganese 

Selenium 

Zinc 

Erbium 

Mercury 

Silicon 

Zirconium 

An  inspection  of  this  table  shows  that  many  elements  are  familiar 
to  us  in  everyday  life.  Thus  copper,  silver,  gold,  iron,  lead,  zinc  and 
tin  are  elements,  as  are  sulphur,  phosphorus,  carbon  and  oxygen. 

Copper,  silver,  gold,  lead  and  a  number  of  other  elements  are  called 
metals;  they  show  metallic  lustre  and  conduct  heat  and  electricity. 
The  elements  which  do  not  possess  these  properties  are  called  non- 
metals.  This  class  includes  oxygen,  hydrogen,  sulphur,  phosphorus, 
carbon,  and  many  other  elements,  the  properties  of  which  are  con- 
sidered in  detail  in  the  earlier  part  of  the  book.  The  division  of  the 
elements  into  metals  and  non-metals  is  not  in  all  respects  a  satis- 
factory one,  but  the  lines  on  which  the  elements  can  be  classified  can 
only  be  adequately  discussed  after  we  have  become  familiar  with  their 
more  important  physical  and  chemical  properties. 

As  regards  the  physical  state  of  the  elements,  a  few  of  them,  such 
as  hydrogen  and  oxygen,  are  gases  at  the  ordinary  temperature,  two 
only,  mercury  and  bromine,  are  liquid,  and  the  great  majority  are 
solid  under  ordinary  conditions. 

A  few  of  the  elements,  such  as  copper,  silver,  gold,  oxygen  and 
sulphur  occur  free  (that  is  as  elements)  in  nature,  but  most  of  them 
are  found  naturally  only  in  the  form  of  chemical  compounds.  The 


THE   CHEMICAL   ELEMENTS 


methods  employed  in  obtaining  elements  from  their  compounds  are 
considered  in  detail  in  connexion  with  the  individual  elements.  We 
have  already  seen  that  heat  and  electrical  energy  can  be  successfully 
employed  for  this  purpose  in  some  cases. 

The  relative  proportions  in  which  the  elements  occur  in  the  part  of 
the  earth  accessible  to  us  are  very  unequal.  Oxygen,  which  occurs 
in  the  air,  in  water,  and  in  the  earth's  crust,  constitutes  about  50  per 
cent,  by  weight  of  the  part  of  the  earth  (including  the  sea)  known  to 
us.  The  nine  elements,  oxygen,  silicon,  aluminium,  iron,  calcium, 
magnesium,  sodium,  potassium,  and  hydrogen  (free  or  in  combination) 
together  make  up  more  than  99  per  cent,  of  the  earth's  crust. 

The  following  table,  compiled  by  Clarke,  contains  an  estimate  of 
the  relative  amounts  of  the  elements  occurring  in  the  part  of  the 
earth  known  to  us  (to  the  depth  of  half  a  mile). 


Earth's  Crust. 

Ocean. 

Atmosphere. 

Entire  Earth. 

Per  cent. 

Per  cent. 

Per  cent. 

Per  cent. 

Oxygen       . 

47.29 

85-79 

23.00 

49.98 

Silicon    . 

27.21 

2  5-  3° 

Aluminium 

7.81 

... 

M 

7.26 

Iron  .     .     . 

5.46 

M 

5-o8 

Calcium 

3-77 

0.05 

3-51 

Magnesium 

2.08 

0.14 

2.50 

Sodium  . 

2.36 

0.14 

M 

2.28 

Potassium  . 

2.40 

0.04 

- 

2.23 

Hydrogen  , 

O.2I 

10.67 

M 

0.94 

Titanium    . 

0.33 

o  30 

Carbon  . 

0.22 

0.002 

0.21 

Chlorine      . 

O.OI 

2.07 

. 

O.IS 

Bromine 

0.008 

. 

Phosphorus 

O.IO 

O.09 

Sulphur.     . 

0.03 

0.09 

0.04 

Nitrogen     . 

... 

77.00 

O.02 

CHAPTER  IV 
OXYGEN— COMBUSTION 

WITH  this  chapter  we  commence  the  systematic  study  of  the 
elements  and  their  more  important  compounds.  In  general, 
the  order  of  treatment  will  be  as  follows :  History,  Occurrence, 
Methods  of  Preparation,  Physical  Properties,  and  Chemical  Properties 
of  the  particular  element.  The  more  important  compounds  which  the 
element  forms  with  other  elements  will  then  be  considered. 

History — The  discovery  of  oxygen  was  made  and  publicly  an- 
nounced by  the  English  chemist  Priestley  in  1774.  He  obtained  it  by 
enclosing  red  oxide  of  mercury  in  a  tube  over  mercury,  and  concentrat- 
ing the  rays  of  the  sun  on  it  by  means  of  a  powerful  lens.  Priestley 
observed  that  substances  which  burned  in  air  burned  still  more  rapidly 
in  the  new  gas,  and,  for  reasons  which  will  be  mentioned  later  (p.  29), 
he  termed  the  gas  dephlogisticated  air.  Oxygen  was  independently 
discovered  by  the  Swedish  chemist  Scheele,  who  in  1775  described 
a  number  of  methods  of  preparing  it.  A  recent  study  of  Scheele's 
original  papers  has  shown,  however,  that  this  chemist  discovered 
oxygen  in  1773,  about  a  year  before  Priestley.  In  the  publication 
of  his  discovery,  Scheele  was,  however,  anticipated  by  Priestley,  as 
already  indicated. 

Lavoisier  observed  that  many  of  the  products  obtained  by  burning 
substances  in  Priestley's  dephlogisticated  air  readily  dissolved  in  water, 
and  the  resulting  solutions  possessed  a  sour  taste  and  showed  certain 
other  properties  generally  regarded  as  characteristic  of  acids  (p.  98). 
For  this  reason  he  called  the  new  gas  oxygen,  or  acid  producer  (o£us-, 
acid,  and  ycvvda),  I  produce),  and  expressed  the  view  that  it  is  the 
acidifying  principle,  to  which  acids  owe  their  characteristic  properties. 
Later  investigation  has,  however,  shown  that  this  view  is  erroneous, 
as  some  of  the  best  known  acids,  for  example  hydrochloric  acid, 
contain  no  oxygen. 

Occurrence — Free  oxygen  is  a  very  important  constituent  of  the 
atmosphere,  in  which  it  occurs  mixed  with  about  four  times  its  volume 
of  another  colourless  gas,  called  nitrogen. 


OXYGEN— COMBUSTION 


21 


In  chemical  combination  it  constitutes  about  eight-ninths  of  the 
weight  of  water,  the  remaining  ninth  being  hydrogen.  It  also  forms 
a  very  important  constituent  of  nearly  all  rocks  and  earthy  substances 
— in  fact,  44  to  48  per  cent,  of  that  part  of  the  crust  of  the  earth  known 
to  us  consists  of  combined  oxygen  (p.  19).  Oxygen  is  also  one  of 
the  essential  constituents  of  practically  all  animal  and  vegetable 
substances. 

Chemical  compounds  in  which  oxygen  is  combined  with  only  one 
other  element  are  termed  oxides. 

Preparation— Laboratory  Methods— (i)  Oxygen  may  be  pre- 
pared in  small  amount  by  heating  mercuric  oxide,  as  already  de- 


FIG.  8. 

scribed  (p.  5).  For  this  purpose  a  small  quantity  of  the  oxide  is 
placed  in  a  hard-glass  tube,  closed  at  one  end,  and  provided  with 
a  cork  and  delivery-tube,  as  shown  in  Fig.  8.  For  the  collection  and 
manipulation  of  gases  over  water  the  arrangement  shown  in  the 
figure  is  very  convenient.  The  hollow  cylindrical  vessel  A,  which 
has  a  small  central  opening  on  the  upper  surface,  and  is  open 
on  the  lower  side,  is  placed  in  the  vessel  B,  and  completely  covered 
by  water.  The  end  of  the  delivery  tube  is  then  placed  below  the 
lower  edge  of  the  vessel  A,  which  is  grooved  for  the  purpose,  and 
the  hard-glass  tube  carefully  heated  with  a  Bunsen  burner.  The  gas 
is  allowed  to  bubble  through  the  water  for  some  time  in  order  to 


22     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

drive  out  the  air.     A  gas-collecting  jar  is  then  inverted  over  A,  as 
shown,  when  the  oxygen  displaces  the  water  in  the  jar. 

If,  as  is  necessary  in  some  cases,  a  gas  has  to  be  collected  over 
mercury,  a  vessel  known  as  a  pneumatic  trough,  provided  with  a 
permanent  shelf  in  the  interior,  is  used  for  the  purpose. 

(2)  Oxygen  can  be  obtained  most  readily  for  laboratory  purposes 
by    heating    potassium    chlorate,    a    substance    which    occurs    in 
colourless  crystals.    When  the  chlorate  is  heated  in  a  test-tube  it 
melts,  and  at  a  rather  higher  temperature  is  slowly  split  up  into 
oxygen  and  another  colourless  substance  termed  potassium  chloride. 
The  reaction  is  one  of  simple   decomposition  (p.  7),  and  is  repre- 
sented by  the  equation  l 

Potassium  Chlorate  =  Potassium  Chloride  +  Oxygen. 

If  the  potassium  chlorate  is  mixed  with  about  one-fourth  of  its 
weight  of  a  black  powder  called  manganese  dioxide,  oxygen  is  given 
off  rapidly  at  a  temperature  below  the  melting-point  of  the  chlorate. 
This  is  the  most  convenient  method  for  the  laboratory  preparation  of 
oxygen,  although  the  gas  obtained  in  this  way  is  not  quite  pure.  The 
gas  may  be  collected  over  water,  as  described  on  the  previous 
page,  and  several  jars  should  be  filled  with  it,  in  order  to  demonstrate 
its  properties  (p.  25). 

The  most  remarkable  fact  about  the  change  just  described  is  that 
although  the  manganese  dioxide  greatly  accelerates  the  liberation  of 
oxygen  from  potassium  chlorate,  and  thus  enables  it  to  take  place  at 
a  much  lower  temperature  than  when  the  chlorate  is  heated  alone, 
yet  the  manganese  dioxide  can  be  recovered  unaltered  in  amount  at 
the  end  of  the  process,  and  therefore  plays  no  apparent  part  in  the 
reaction.  This  is  only  one  of  many  instances  in  which  a  substance 
greatly  accelerates  a  chemical  change,  while  itself  remaining  unaltered 
at  the  end  of  the  reaction.  In  such  cases  the  substance  is  said  to 
exert  a  catalytic  effect,  and  is  termed  a  catalyst  for  the  change  in 
question.  It  should  be  clearly  understood  that  the  use  of  these  terms 
does  not  suggest  any  explanation  for  the  effect  in  question,  but  merely 
classes  together  a  number  of  phenomena  having  certain  important 
features  in  common.  We  shall  meet  with  many  examples  of  catalytic 
action  in  the  course  of  our  subsequent  work. 

(3)  Very   pure   oxygen   may  be  obtained    by  heating  potassium 
permanganate,  a  substance  which   occurs  in  reddish-black  crystals. 

1  The  term  equation  denotes  that,  in  accordance  with  the  law  of  conservation  of 
weight,  the  sum  of  the  weights  of  the  substances  on  one  side  of  the  =  sign  is 
equal  to  that  on  the  other  (cf.  p.  10). 


OXYGEN— COMBUSTION  23 

The  experiment  is  carried  out  as  described  under  (i).  The  chemical 
change  in  this  case  is  rather  complex,  and  will  be  considered  under 
potassium  permanganate  (p.  548). 

(4)  Oxygen   may  be   obtained   by  the    effect    of   heat   on   many 
substances    other  than   those   mentioned,   for  example,  manganese 
peroxide,  but   higher  temperatures  are   usually   required,  and   the 
reactions  are  in  other  respects  less  suitable  for  laboratory  purposes 
than  those  already  described. 

(5)  Oxygen  may  also  be  obtained  from  substances  such  as  potassium 
permanganate  and  potassium  bichromate  by  heating  with  sulphuric 
acid.     These  methods  are  dealt  with  in  detail  at  a  later  stage. 

Commercial  Preparation  of  Oxygen— (6)  Brirts  Oxygen 
Process — Oxygen  gas,  compressed  into  steel  cylinders,  is  now  an 
article  of  considerable  commercial  importance.  For  preparing  a 
small  quantity  of  oxygen  in  the  laboratory,  the  convenience  of  the 
method  is  the  chief  consideration,  and  the  cost  a  matter  of  secondary 
importance.  In  preparing  a  substance  on  the  commercial  scale, 
however,  the  cost  is  of  prime  importance,  and  it  is  therefore  necessary 
to  find  a  plentiful  source  of  the  raw  material  from  which  the  substance 
required  may  readily  be  obtained.  In  the  case  of  oxygen,  it  is 
natural  to  think  of  the  atmosphere  as  a  source  of  supply,  since  it 
is  mainly  composed  of  free  oxygen  and  nitrogen.  A  method  suitable 
for  the  present  purpose  would  be  to  use  a  substance  which  enters  into 
chemical  combination  with  oxygen,  but  not  with  nitrogen,  and  the 
resulting  compound  must  readily  give  up  its  oxygen  under  suitable 
conditions. 

A  substance  which  answers  these  requirements  is  barium  oxide. 
When  this  substance,  in  the  form  of  a  white  powder,  is  heated  at 
low  red  heat  (about  500°)  in  the  air,  it  takes  up  oxygen  forming 
a  compound  called  barium  peroxide,  which,  for  the  same  proportion 
of  barium,  contains  twice  as  much  oxygen  as  the  oxide.  When  the 
peroxide  is  heated  to  a  bright  red  heat  (about  1000°),  it  gives  up  half 
of  its  oxygen,  and  barium  oxide  is  reformed.  The  same  quantity  of 
barium  oxide  may  be  used  over  and  over  again,  being  alternately 
heated  while  air  is  passed  over  it,  and  the  product  then  raised  to  ft 
higher  temperature  to  drive  off  part  of  the  oxygen. 

The  reaction  may  conveniently  be  represented  by  the  following 
equation 

Barium  peroxide^tBarium  oxide  +  oxygen 

in  which  the  change  in  the  direction  of  the  upper  arrow  takes  place 
at  a  bright  red  heat ;  that  in  the  direction  of  the  lower  arrow  at  a 


24     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

low  red  heat.    A  change  of  this  type,  which  may  proceed  in  either 
direction  depending  on  the  conditions,  is  termed  a  reversible  reaction. 

In  actual  practice,  the  decomposition  of  the  barium  peroxide  is 
effected  more  economically  by  reducing  the  pressure  instead  of  by 
raising  the  temperature.  The  process  thus  consists  in  passing  air 
over  heated  barium  oxide  at  atmospheric  pressure,  and  then  removing 
the  oxygen  at  the  same  temperature  under  reduced  pressure.  Before 
being  admitted  to  the  barium  oxide,  the  air  is  passed  over  lime  to 
remove  a  gas  called  carbon  dioxide,  which  otherwise  would  combine 
with  the  barium  oxide  and  render  it  useless  for  absorbing  oxygen. 

The  process  just  described  is  termed  Brin's  oxygen  process.  The 
compressed  oxygen  is  not  pure,  but  contains  a  few  per  cent,  of 
nitrogen. 

(7)  On  the  commercial  scale,  oxygen  is  now  obtained  almost 
entirely  by  the  evaporation  of  liquid  air.  The  latter  is  a  mixture  of 
liquid  oxygen  and  nitrogen.  As  liquid  nitrogen  goes  into  vapour 
more  readily  than  does  liquid  oxygen,  a  mixture  rich  in  oxygen  can 
be  obtained  when  liquid  air  is  allowed  to  evaporate  slowly.  As  a 
number  of  factors  are  concerned  in  the  process,  which  at  this  early 
stage  have  not  yet  been  met  with,  the  consideration  of  this  method  of 
preparing  oxygen  is  postponed  to  the  section  dealing  with  the  lique- 
faction of  gases  (p.  73). 

Physical  Properties — Oxygen  is  a  colourless,  odourless,  taste- 
less gas.  It  is  1.105  times  heavier  than  air.  The  usual  standard  to 
which  the  density  of  gases  is  referred  is  that  of  hydrogen  (the  lightest 
gas)  taken  as  unity;  on  this  basis  the  density  of  oxygen  is  15.90 
The  weight  of  a  litre  of  oxygen  at  o°  and  760  mm.  pressure  is, 
according  to  Morley,  1.4290  grams,  according  to  Rayleigh  1.4295 
grams. 

Oxygen  may  be  liquefied  at  and  below  —  119°,  which  is  its  critical 
temperature  (p.  71).  At  its  critical  temperature  about  58  atmos- 
pheres pressure  are  required  to  liquefy  it  (the  so-called  critical 
pressure),  and  the  further  the  temperature  is  lowered  below  the 
critical  temperature  the  less  is  the  pressure  required  to  convert  the 
gas  to  a  liquid.  At  its  boiling-point  the  specific  gravity  of  the  liquid 
is  1.131.  Devvar  has  succeeded  in  obtaining  oxygen  as  a  bluish, 
snow-like  solid  by  cooling  with  liquid  hydrogen  ;  the  solid  melts  at 
—  227°.  Liquid  oxygen  is  attracted  by  a  magnet. 

Oxygen  is  slightly  soluble  in  water  ;  I  c.c.  of  water  dissolves  at  o° 
0.0489  c.c.,  at  20°  0.031  c.c.,  and  at  30°  0.026  c.c.  of  the  gas,  measured 
in  each  case  under  i  atmosphere  pressure  (760  mm.).  According  to 


OXYGEN— COMBUSTION 


Winkler,  the   solubility,  C,  of  oxygen  in   water  diminishes  as  the 
temperature  rises  in  accordance  with  the  formula 

0  =  0.0489  -  0.001341 3/+0.0000283/2- 0.00000029534/8. 

It  is  a  remarkable  fact  that  oxygen  is  fairly  soluble  in  fused  silver, 
but  escapes  almost  entirely  when  the  silver  solidifies. 

Chemical  Properties— In  describing  the  chemical  properties 
of  an  element,  we  are  mainly  concerned  with  its  power  of  combining 
directly  with  other  elements,  the  conditions  under  which  the  chemical 
changes  occur,  and  the  nature  of  the  compounds  formed.  When  an 
element  has  the  power  of  combining  readily  with  a  large 
number  of  other  elements  it  is  said  to  be  chemically 
active. 

Many  substances,  such  as  sulphur,  phosphorus,  carbon, 
and  iron,  combine  rapidly  with  oxygen  when  the  reaction 
has  been  started  by  heating.  This  may  readily  be  shown 
as  follows.  A  number  of  gas  jars  are  rilled  with  oxygen 
by  heating  a  mixture  of  potassium  chlorate  and  man- 
ganese dioxide  (p.  22),  and  are  closed  by  glass  covers 
(lubricated  with  vaseline  or  grease)  until  required.  The 
sulphur  (phosphorus  or  carbon)  is  placed  in  a  spoon  of 
the  form  shown  in  Fig.  9,  heated  till  it  just  begins  to  burn, 
and  then  plunged  into  the  oxygen,  when  it  continues  to 
burn,  but  much  more  rapidly  than  in  air.  In  order  to 
show  the  burning  of  iron  in  oxygen,  a  bundle  of  thin  iron 
wires  is  bent  round  a  small  piece  of  sulphur  at  one  end, 
and  this  end  is  then  heated  in  a  flame.  The  combination 
of  the  iron  and  sulphur  gives  out  a  considerable  amount 
of  heat,  which  starts  the  combination  of  the  iron  and 
oxygen  of  the  air.  If  the  wires,  held  in  a  tongs,  are  quickly  immersed 
in  a  jar  of  oxygen,  they  continue  to  burn  with  an  extremely  brilliant 
flame,  giving  rise  to  a  shower  of  sparks,  while  particles  of  the  fused 
product  of  the  reaction  fall  on  the  bottom  of  the  gas  jar  and  solidify. 

If,  after  action  has  ceased,  the  contents  of  the  jars  are  shaken  up 
with  water,  and  a  few  drops  of  litmus  solution  added,  it  will  be 
noticed  that  the  litmus  is  turned  red  in  the  jars  in  which  sulphur, 
phosphorus,  and  carbon  have  been  burned,  whilst  no  change  of  colour 
occurs  in  the  jar  in  which  iron  has  been  burned.  The  change  in  the 
colour  of  the  litmus  to  red  indicates  the  presence  of  an  acid  (p.  98), 
from  which  it  follows  that  the  oxides  obtained  by  burning  sulphur, 
phosphorus,  and  carbon  in  the  air  dissolve  in  water  to  form  acids. 


FIG.  9. 


26     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Oxides  which  dissolve  in  water  to  form  acids  are  termed  acidic  oxide* 
or  anhydrides. 

The  products  formed  when  substances  burn  in  air  are  the  same  as 
those  formed  when  they  burn  in  oxygen,  the  only  difference  being  that 
in  the  latter  case  the  reactions  are  more  vigorous.  This  is  easily 
understood  when  it  is  remembered  that  the  atmosphere  is  essentially 
a  mixture  of  oxygen  and  another  gas,  nitrogen;  as  the  latter  is 
a  comparatively  indifferent  gas,  and  combines  directly  with  very  few 
substances,  air  behaves  towards  the  great  majority  of  elements  simply 
as  diluted  oxygen. 

Oxygen  combines  with  all  the  other  elements  except  fluorine  and 
the  rare  elements  helium,  neon,  argon,  krypton,  and  xenon.  In  the 
case  of  a  few  elements,  such  as  silver  and  gold,  the  oxides  can  only 
be  obtained  indirectly. 

When  an  element  or  compound  enters  into  chemical  combination 
with  oxygen,  the  process  is  termed  oxidation,  and  the  element  or 
compound  is  said  to  be  oxidizea. 

Combustion — The  rapid  combination  of  certain  elements  with 
oxygen  forms  a  good  illustration  of  combustion,  which  may  be  defined 
as  a  chemical  change  which  proceeds  with  the  evolution  of  light  and 
heat.  From  the  definition  it  follows  that  the  term  combustion  is  not 
confined  to  combinations  with  oxygen,  although  these  are  by  far  the 
most  familiar.  When  iron  filings  and  sulphur  are  heated  together 
in  a  test-tube,  they  combine  with  evolution  of  light  and  heat,  and 
therefore,  according  to  the  above  definition,  this  is  a  process  oi 
combustion. 

The  combination  of  many  elements  with  oxygen,  which  proceeds 
very  rapidly  at  high  temperatures,  may  also  proceed  at  a  low  tem- 
perature, though  much  more  slowly.  The  familiar  glowing  of  phos- 
phorus in  the  dark  is  an  accompaniment  of  the  combination  of  oxygen 
and  phosphorus,  and  leads  to  the  formation  of  the  same  compound 
as  when  phosphorus  burns  brightly.  Further,  the  amount  of  energy 
given  out  when  a  definite  amount  of  an  element  combines  with 
oxygen  is  the  same  whether  the  action  takes  place  quickly  or  slowly. 
When  combustion  is  rapid,  the  heat  is  given  out  quickly  and  can 
thus  raise  the  products  of  combustion  to  a  high  temperature.  When, 
however,  the  change  is  slow,  the  heat  of  reaction  escapes  into  the 
surroundings  and  the  temperature  may  not  rise  much  above  the 
ordinary  temperature. 

Importance  of  the  Study  of  Combustion  for  the  De- 
velopment of  Chemistry — Many  of  the  chemical  phenomena 


OXYGEN— COMBUSTION  27 

with  which  we  are  familiar  in  everyday  life,  such  as  the  burning  oi 
coal  and  wood  in  air,  the  rusting  of  iron,  the  glowing  of  phosphorus, 
the  burning  of  sulphur  in  air,  are  processes  of  combustion.  A 
point  of  fundamental  importance  in  all  these  processes  is  that  the 
product  or  products  of  combustion  weigh  more  than  the  original 
substance.  At  first  sight  there  appears  to  be  a  loss  of  weight  in  some 
of  these  reactions,  for  example,  when  a  candle  burns  in  air ;  but  it  has 
already  been  pointed  out  that  when  the  experiment  is  performed  in 
such  a  way  that  the  products  of  combustion  are  not  allowed  to  escape 
into  the  atmosphere,  there  is  an  evident  increase  of  weight  in  this 
case  also. 

When  it  is  remembered  that  oxygen  is  one  of  the  constituents  of 
air,  these  observations  are  readily  understood.  The  burning  of  the 
candle  is  essentially  a  process  of  chemical  combination  between  its 
constituents  and  oxygen,  and  it  may  be  anticipated  that,  in  accordance 
with  the  law  of  conservation  of  mass,  the  excess  in  weight  of  the 
products  is  just  balanced  by  the  loss  in  weight  of  the  atmosphere. 
Although  this  explanation  of  the  phenomena  of  combustion  is  simple 
and  easily  understood,  it  was  only  arrived  at  after  much  experiment 
and  discussion.  Its  general  acceptance  towards  the  end  of  the 
eighteenth  century  was  due  to  Lavoisier.  A  brief  account  of  his 
classical  experiment  on  the  oxidation  of  mercury  in  a  confined  volume 
of  air,  and  of  the  conclusions  he  drew  from  his  observations,  will  now 
be  given. 

Four  ounces  of  mercury  were  placed  in  the  glass  retort  A  (Fig.  10), 
the  neck  of  which  was  bent  as  shown  and  dipped  into  the  glass  vessel 
B,  which  contained  air  confined  over  mercury.  The  retort  containing 
the  mercury  was  then  heated  for  twelve  days  at  a  temperature  near 
the  boiling-point  of  mercury,  and  it  was  noticed  that  the  surface  of 
the  metal  gradually  became  covered  with  red  scales  (mercuric  oxide). 
After  the  apparatus  had  cooled,  it  was  noticed  that  the  volume  of  the 
confined  air  had  diminished  ;  out  of  50  cubic  inches  originally  taken 
only  42  remained.  The  red  scales  were  collected,  and  on  heating 
gave  off  7-8  cubic  inches  of  a  gas  (oxygen)  which  supported  com- 
bustion much  more  energetically  than  ordinary  air.  Lavoisier  also 
found  that  the  gas  remaining  in  the  vessel,  which  had  not  combined 
with  the  mercury,  was  no  longer  capable  of  oxidizing  metals,  and 
a  lighted  taper  was  extinguished  in  the  gas  as  if  it  had  been  plunged 
into  water.  It  follows  from  this  experiment  that  air  is  a  mixture 
of  two  gases,  one  of  which  (oxygen)  combines  with  metals  such  as 
mercury,  the  other,  present  in  larger  proportion  (nitrogen),  does  not 


28     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

combine  with  metals  and  does  not  support  the  combustion  of  a  lighted 
taper. 

The  further  point,  that  the  gain  in  weight  of  the  products  is  balanced 
by  a  loss  in  weight  of  the  atmosphere,  was  proved  more  conclusively 
by  Lavoisier  in  a  further  experiment  on  the  combustion  of  tin.  Some 
pieces  of  tin  were  sealed  up  in  a  glass  vessel  along  with  air,  and 
the  vessel  heated  for  some  time.  It  was  then  found  that,  although 
the  tin  had  altered  in  appearance  and,  as  subsequent  investigation 
showed,  had  gained  in  weight,  there  was  no  difference  in  the  weight 
of  the  closed  vessel  before  and  after  heating.  When,  however,  the 
vessel  was  cautiously  opened  air  rushed  in,  and  it  was  found  that 


FIG.  10. 

the  weight  of  the  air  which  entered  was  equal  to  the  gain  in  weight 
of  the  tin  during  the  heating. 

The  simplest  laboratory  method  of  showing  that  a  gas  is  taken  up 
from  the  atmosphere  in  the  process  of  combustion  is  to  burn  a  piece 
of  phosphorus  in  a  confined  volume  of  air.  The  phosphorus  is 
placed  on  an  inverted  crucible  lid  floating  on  water,  and  the  whole 
is  covered  by  a  bell-jar  provided  with  a  well-fitting  stopper  (Fig.  1 1). 
The  stopper  is  removed  for  a  moment  and  the  water  brought  to 
the  same  level  inside  and  outside  the  jar,  the  phosphorus  ignited 
by  touching  with  a  heated  rod,  the  stopper  immediately  replaced, 
and  the  bell-jar  pressed  down  against  some  blotting-paper  or 
other  soft  material  in  the  bottom  of  the  dish  while  the  combustion 


OXYGEN— COMBUSTION 


29 


lasts.1  The  bell-jar  becomes  filled  with  white  fumes  (of  phosphorus 
pentoxide),  and  when  practically  all  the  oxygen  is  used  up  the  phos- 
phorus becomes  extinguished.  The  jar  is  then  cautiously  raised 
without  allowing  the  lower  edge  to  rise  above  the  surface  of  the 
water,  when  it  will  be  found  that  water  enters  and  fills  about  one-fifth 


FIG.  ii. 

of  the  space  previously  occupied  by  air.  The  gas  remaining  in  the 
jar  is  incapable  of  supporting  the  combustion  of  a  burning  taper. 
The  conclusion  which  might  be  drawn  from  this  experiment,  that 
about  one-fifth  of  the  atmosphere  by  volume  is  oxygen,  is  fully 
confirmed  by  further  investigations  (p.  203)  ;  the  remaining  four-fifths 
is  mainly  nitrogen. 

The  Phlogiston  Theory.— Before  Lavoisier's  time  another  view  of  the  nature 
of  combustion,  termed  the  phlogiston  theory,  was  held  almost  universally  through- 
out the  chemical  world.  According  to  this  view  combustion  is  essentially  a  process 
of  decomposition.  Every  combustible  substance  contains  a  certain  proportion  of 
a  common  constituent  termed  by  Stahl  phlogiston ,  and  in  the  process  of  combus- 
tion splits  up  into  phlogiston  and  another  product.  Thus  sulphur  was  supposed 
to  consist  of  sulphuric  acid  and  phlogiston,  and  to  decompose  into  those  two  sub- 
stances on  combustion ;  a  metal,  such  as  lead,  was  supposed  to  be  composed  of 
its  calx  (or,  as  we  would  say,  its  oxide)  and  phlogiston,  and  so  on.  Phlogiston 
itself  was  not  of  course  known  in  the  free  condition,  but  combustible  substances 

1  The  air  expands  at  first  owing  to  the  heat  given  out  in  the  combustion,  and 
part  of  it  may  escape  below  the  edge  of  the  bell-jar  unless  the  precaution  above 
mentioned  is  taken. 


30     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

contained  it  to  a  greater  or  less  extent,  and  carbon  was  supposed  to  be  nearly 
pure  phlogiston.  It  was  therefore  easy  to  understand  how  a  substance  deprived 
of  its  phlogiston,  such  as  the  calx  of  a  metal,  could  be  restored  to  its  original 
condition  by  heating  with  carbon.  The  latter,  containing  a  large  proportion  of 
phlogiston,  gives  up  this  material  to  the  calx,  and  the  metal,  rich  in  phlogiston, 
is  regained.  After  the  discovery  of  hydrogen,  and  the  observation  that  it  could 
change  calxes  to  metals,  this  substance  was  regarded  as  pure  phlogiston. 

The  chief  difficulty  of  the  phlogiston  theory,  which  ultimately  led  to  its  over- 
throw, was  the  fact,  established  long  before  Stahl's  day,  that  combustion  is 
attended  by  a  gain  in  weight,  whereas  if  phlogiston  is  a  material  substance  there 
ought  to  be  a  loss  of  weight  during  combustion.  The  adherents  of  the  phlogiston 
theory  made  many  ingenious  attempts  to  overcome  this  difficulty,  one  suggestion 
being  that  phlogiston  had  a  negative  weight,  and  therefore  on  its  escape  the 
residue  weighed  more  than  before.  It  seems  a  little  curious  that  so  able  an 
investigator  as  Boyle,  who  knew  from  his  own  experiments  that  combustion  is 
attended  by  an  increase  in  weight,  that  air  is  necessary  for  combustion  and  that 
part  of  it  is  absorbed  in  the  process,  should  have  been  doubtful  whether  lead  calx 
is  a  constituent  of  lead  or  lead  a  constituent  of  lead  calx  (lead  oxide). 

It  should,  however,  be  recognized  that  the  phlogiston  theory,  although  unsuit- 
able in  many  respects,  contributed  materially  to  the  development  of  chemistry.  It 
will  be  clear  from  what  has  been  mentioned  above  with  regard  to  the  conversion 
of  lead  to  its  calx  and  the  reconversion  of  the  latter  to  metallic  lead  by  charcoal, 
that  the  escape  of  phlogiston  denotes  what  we  now  term  oxidation,  and  the 
addition  of  phlogiston  constitutes  reduction.  On  this  basis  the  chemists  of  that 
period  were  enabled  to  classify  together  many  changes  previously  regarded  as 
quite  different,  which  greatly  facilitated  experimental  investigation. 

The  function  of  the  air  in  combustion,  according  to  the  adherents  of  the  phlogiston 
theory,  was  to  take  up  phlogiston,  and  when  it  was  no  longer  able  to  support 
combustion  it  was  said  to  be  saturated  with  phlogiston.  As  oxygen  was  a  much 
better  supporter  of  combustion  than  ordinary  air,  it  is  now  easy  to  see  why  it 
was  termed  by  Priestley  dephJogisticated  air.  On  the  same  principle,  nitrogen 
was  termed  by  Priestley  phlogisticated  air. 


CHAPTER  V 
HYDROGEN— GENERAL  PROPERTIES  OF  GASES 

History — Paracelsus,  in  the  sixteenth  century,  was  familiar  with 
the  fact  that  an  inflammable  gas  is  produced  by  the  action 
of  dilute  acids  on  certain  metals,  but  the  gas  was  then  looked  upon 
as  a  form  of  ordinary  air.  Cavendish,  in  1766,  was  the  first  to 
prepare  pure  hydrogen  and  to  recognize  it  as  a  definite  substance ; 
he  obtained  it,  like  Paracelsus,  by  the  action  of  acids  on  metals. 
As  has  already  been  pointed  out,  hydrogen  is  a  constituent  of  water, 
hence  its  name,  which  is  derived  from  the  Greek  words  v6\ap,  water, 
and  y«Wo>,  I  produce. 

Occurrence — Free  hydrogen,  mixed  with  other  gases,  is  given 
off  from  volcanoes,1  and  is  also  found  in  small  amount  enclosed  in 
meteorites.  It  also  occurs  free  in  traces  in  the  atmosphere  ;  accord- 
ing to  Rayleigh  the  average  proportion  does  not  exceed  I  in  30,000 
by  volume,  whilst  according  to  Ramsay  the  amount  does  not  exceed 
i  in  a  million.  Spectroscopic  observations  indicate  that  free  hydrogen 
is  present  in  very  large  proportion  in  the  sun,  and  in  many  fixed  stars 
and  nebulae. 

In  the  combined  form,  hydrogen  occurs  very  largely  on  the  earth. 
It  forms  over  11  per  cent,  of  water,  and  is  one  of  the  essential  ele- 
ments in  plants  and  animals.  It  also  occurs,  combined  with  carbon, 
in  natural  gas,  marsh  gas,  and  petroleum.  It  is  an  essential  con- 
stituent of  acids. 

Preparation — (A)  From  Water.  As  water  is  a  chemical  com- 
pound of  hydrogen  and  oxygen,  and  is  always  available,  it  is  natural 
to  have  recourse  to  it  as  a  source  of  hydrogen.  The  elements  may 
be  considered  as  being  held  together  by  chemical  attraction,  and  this 
attraction  must  be  overcome  in  some  way  in  order  to  obtain  free 
hydrogen.  Another  and  preferable  way  of  regarding  the  matter 
is  that  hydrogen  and  oxygen  give  out  a  large  amount  of  energy, 
mainly  in  the  form  of  heat,  when  they  combine  to  form  water,  and 

1  The  gases  given  off  from  Mount  Pel^e,  Martinique,  during  the  eruption  of 
1902,  contained  over  20  per  cent,  by  volume  of  free  hydrogen  (Moissan). 

31 


32     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

in  order  to  obtain  the  free  elements  from  water  this  energy  must  be 
supplied  in  some  way  (p.  14).  From  these  considerations  two 
principal  methods  might  be  suggested  for  obtaining  hydrogen  from 
water :  (a)  to  supply  a  large  amount  of  energy  to  water,  for  example, 
in  the  form  of  heat  or  of  electrical  energy ;  (b)  to  bring  water  in  contact 
with  a  substance  which  has  a  greater  attraction  for  oxygen  than  the 
oxygen  has  for  hydrogen.  The  more  important  methods  for  the 
preparation  of  hydrogen  from  water,  and  incidentally  for  the  de- 
composition of  water,  will  now  be  briefly  considered. 

(1)  When  steam  is  passed  through  a  platinum  tube,  heated  to  a 
temperature  exceeding   1000°,  and  the  gases  leaving  the  tube  are 
rapidly  cooled,  very  small  amounts   of  hydrogen  and  oxygen  are 
obtained.     This  method  is  not,  of  course,  a  practical  one  for  pre- 
paring  hydrogen    in    quantity,  but    is    of   considerable    theoretical 
interest.     It  is  further  referred  to  in  connexion  with  water  (p.  52). 

(2)  Hydrogen,  and   also  oxygen,   can   readily  be  obtained   from 
water  by  the  employment  of  electrical  energy.     The  apparatus  used 
for  this  purpose,  known  as  a  voltameter,  and  the  method  of  pro- 
cedure,  have  been   described   in   a  previous    chapter.      Hydrogen 
obtained  in  this  way  is  very  pure. 

(3)  The  second  general  method  of  obtaining  hydrogen  from  water 
can  be  illustrated  by  means  of  the  soft  metals  sodium  and  potassium, 
which  are  acted  on   by  water    at   the   ordinary   temperature   with 
liberation  of  hydrogen.     With  the  former  metal  the  experiment  is 
best  performed  by  enclosing  a  small   piece   in  a   so-called   sodium 
spoon   (a  small  chamber  of  wire-gauze  at  the  end  of  a  long  rod), 
and  holding  it  below  the  surface  of  water  beneath  an  inverted  test- 
tube  filled  with  water.     The  sodium  acts  rapidly  on  water,  liberat- 
ing  hydrogen,  which  displaces    the  water  in   the  test-tube.      The 
experiment  with  potassium  is  best  performed  by  throwing  a  small 
piece  of  the  metal  on  the  surface  of  water  in  a  glass  dish.     Potas- 
sium acts  on   water  more  vigorously  than   sodium   does,  and   the 
heat    given   out  is  so    great   that    the  hydrogen  catches   fire,  and 
burns  with   a  violet  flame,  the   colour  being  due  to  the  vapour  of 
potassium  (p.  417). 

The  water  in  which  sodium  (or  potassium)  has  been  dissolved 
makes  the  fingers  slippery  and  turns  red  litmus  paper  blue.  This 
indicates  the  presence  of  a  new  substance,  which  can  be  obtained 
as  a  white  residue  by  evaporating  the  water.  It  is  called  sodium 
hydroxide  or  caustic  soda.  A  hydroxide,  as  its  name  indicates, 
contains  both  hydrogen  and  oxygen.  The  chemical  change  which 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     33 

occurs  when  sodium  is  thrown  into  water  is,  therefore,  represented 
by  the  equation 

Sodium  +  water  =  sodium  hydroxide  +  hydrogen. 

Substances  which  in  solution  have  a  soapy  feel,  turn  litmus  blue 
and  show  certain  other  characteristic  properties  are  termed  bases, 
and  the  solu  ions  are  said  to  have  an  alkaline  reaction.  Potassium 
hydroxide  and  sodium  hydroxide  are  typical  bases. 

Certain  other  metals,  including  lithium,  barium,  and  calcium,  also 


FIG.  12. 

act  on  water  at  the  ordinary  temperature,  liberating  hydrogen  and 
forming  the  corresponding  hydroxides. 

(4)  Metals  such  as  magnesium  and  zinc,  although  practically  without 
action  on  water  at  room  temperature,  slowly  decompose  boiling  water, 
and  act  still  more  energetically  when  heated  in  a  current  of  steam- 
The  arrangement  in  the  case  of  magnesium  is  illustrated  in  Fig.  12. 
The  metal  is  strongly  heated  in  a  bulb  tube  A  and  steam  from  a 
boiler  B  passed  over  it.  When  nearly  red-hot,  the  metal  catches 
fire  in  the  steam,  forming  magnesium  oxide  and  hydrogen.  The 
latter  can  be  ignited  at  the  end  of  the  tube  as  it  escapes.  At  the 
end  of  the  experiment  the  magnesium  oxide,  in  the  form  of  a  white 
powder,  is  found  in  the  bulb. 
3 


34     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

(5)  Iron  at  a  red  heat  also  decomposes  steam,  forming  hydrogen 
and  an  oxide  of  iron.  The  arrangement  of  the  apparatus  for  this 
purpose  will  readily  be  understood  from  the  illustration  (Fig.  13). 


FIG.  14. 

The  iron  tube,  which  contains  iron  in  the  form  of  nails,  is  raised  to  a 
red  heat  by  means  of  the  burners,  and  steam  is  passed  through  the 
tube  till  all  the  air  is  expelled.  The  hydrogen  escaping  at  the  end  of 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     35 

the  delivery  tube  can  then  be  collected  by  displacement  of  water  and 
tested  in  the  usual  way. 

(6)  The  most  convenient  laboratory  method  for  the  preparation  of 
hydrogen  is  by  the  action  of  zinc  on  dilute  sulphuric  acid.  The  zinc 
is  placed  in  a  two-necked  bottle  (the  so-called  Woulf 's  bottle) ;  one 
opening  is  fitted  with  a  thistle  funnel,  the  other  with  a  gas  delivery 
tube  (Fig.  14).  Dilute  sulphuric  acid  is  poured  into  the  funnel,  and 
the  hydrogen  is  collected  in  the  usual  way. 

It  is  a  remarkable  fact,  which  has  not  yet  been  adequately  ex- 
plained, that  pure  dilute  sulphuric  acid  has  practically  no  action 
on  perfectly  pure  zinc.  If,  however,  a  few  drops  of  a  solution  of 
copper  sulphate  or  of  platinum  chloride  are  added,  the  action  at 
once  becomes  vigorous.  Commercial  granulated  zinc,  which  contains 
other  metals  as  impurities,  is  readily  acted  on  by  sulphuric  acid. 

The  other  product  formed  when  sulphuric  acid  acts  on  zinc  is 
termed  zinc  sulphate,  and  may  be  obtained  in  crystals  on  evaporating 
the  solution.  The  equation  expressing  the  action  of  sulphuric  acid 
on  zinc  is,  therefore,  as  follows  : — 

Zinc  +  sulphuric  acid  =  zinc  sulphate  +  hydrogen. 

Hydrogen  may  also  be  obtained  by  the  action  of  other  acids  on 
metals  other  than  zinc,  as  will  be  shown 
at  a  later  stage. 

In  order  to  obtain  hydrogen  by  the 
action  of  sulphuric  acid  on  zinc  as  re- 
quired, the  apparatus  devised  by  Kipp 
and  represented  in  Fig.  15  is  very  con- 
venient and  economical.  It  consists 
essentially  of  two  glass  bulbs  B  and  C, 
joined  by  a  narrow  glass  tube  ;  into  the 
top  of  the  bulb  B  a  third  bulb  A,  extended 
at  its  lower  part  into  a  long  tube,  is  fitted 
airtight  in  such  a  way  that  the  tube  extends 
nearly  to  the  bottom  of  the  lower  bulb  C. 
The  zinc  is  placed  in  B,  and  is  prevented 
from  falling  down  into  C  by  the  narrow 
neck,  which  is  nearly,  but  not  entirely, 
filled  by  the  prolongation  of  the  bulb  A. 
With  the  stop-cock,  D,  open,  dilute  sul- 


FIG.  15. 


phuric  acid  is  poured  through  A  into  C  till  it  reaches  the  zinc  in  B, 
when  hydrogen  is  given  off.      On  closing  the  stopcock,  gas  continues 


36     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

to  be  given  off  till  all  the  liquid  is  forced  out  of  B  (part  of  it  rising  into 
A  through  the  long  inner  tube),  when  the  action  ceases  automati- 
cally. If  desired,  more  dilute  acid  can  be  poured  into  A  when 
the  stop-cock  D  is  closed,  and  on  opening  D  the  gas  will  be  given 
off  under  increased  pressure,  owing  to  the  higher  level  of  the  liquid 
in  A. 

This  apparatus  may  also  be  used  for  preparing  hydrogen  sulphide, 
carbon  dioxide,  and  other  gases. 

(7)  Although  granulated  zinc  has  very  little  action  on  water  even 
at  boiling-point,  hydrogen  is  readily  given  off  when  water  is  heated 
with  zinc  coated  with  copper,  the  so-called  zinc-copper  couple.  The 
couple  is  prepared  by  immersing  granulated  zinc  for  a  few  minutes 
in  a  dilute  solution  of  copper  sulphate ;  the  excess  of  the  sulphate 
solution  is  then  poured  off,  and  the  couple  washed  two  or  three 
times  with  cold  water.  The  chemical  change  taking  place  when 
the  couple  is  heated  with  water  is  represented  by  the  equation 

Zinc  +  water = zinc  oxide  +  hydrogen, 

so  that  the  copper  apparently  plays  no  direct  part  in  the  change. 

(8)  Very  pure  hydrogen  may  be  obtained  by  heating  metallic  zinc 
with  a  solution  of  sodium  hydroxide.     The  reaction  in  this  case  is 
represented  by  the  equation 

Zinc  +  sodium  hydroxide = sodium  zincate-f-  hydrogen. 
Aluminium  may  be  used  in  place  of  zinc. 

(9)  On  the  commercial  sca/e,  hydrogen  is  obtained  by  electrolysis  of 
sodium  hydroxide  solution  using  iron  electrodes  (p.  15),  or  from  water 
gas  (p.  333)  by  removing  the  carbon  monoxide  by  liquefaction  or  by 
a  chemical  method. 

Physical  Properties— Hydrogen,  like  oxygen,  is  a  colourless, 
odourless,  tasteless  gas.  It  is  the  lightest  gas  known,  its  density 
being  about  one-sixteenth  that  of  oxygen.  One  litre  of  hydrogen,  at 
o°  and  760  mm.  pressure,  weighs  0.089873  grams  (Morley).  The 
lightness  of  hydrogen  may  be  shown  very  strikingly  by  suspending 
a  beaker,  mouth  downwards,  from  one  arm  of  a  balance  and  placing 
weights  in  the  other  pan  till  the  pointer  is  at  zero  on  the  scale.  If 
then  the  contents  of  a  jar  of  hydrogen  are  poured  upwards  into  the 
beaker,  the  movement  of  the  pointer  will  indicate  that  the  beaker 
and  contents  weigh  less  than  before.  On  account  of  its  lightness 
hydrogen  is  employed  for  inflating  balloons. 

Hydrogen  was  first  obtained  as  a  coherent,  colourless,  transparent 
liquid  by  Dewar  in  1898,  and  somewhat  later  by  Travers.  Its  boiling- 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     37 

point  on  the  helium  scale  is  -252.5°,  and  its  melting-point  -259° 
at  a  pressure  of  49-50  mm.  (Travers).  The  critical  temperature  of 
liquid  hydrogen  is  about  -241°  and  its  critical  pressure  14  atmos- 
pheres (Olszewski).  The  density  of  liquid  hydrogen  is  0,07  :  in 
other  words  its  density  is  only  one-fourteenth  that  of  water. 

Hydrogen  is  only  very  slightly  soluble  in  water.  I  c.c.  of  water 
dissolves  at  o°  0.0215  c.c.,  at  10°  0.0198  c.c.,  and  at  20°  0.0184  c«c-  °f 
hydrogen  under  I  atmosphere  pressure  (Timofejeff). 

Chemical  Properties— If  a  jar  is  filled  with  pure  dry  hydrogen, 
and  a  lighted  taper  is  applied  to  the  mouth  of  the  jar,  the  gas  catches 
fire  and  burns  with  an  almost  colourless  flame.  After  the  flame  has 
gone  out,  moisture  will  be  observed  on  the  sides  of  the  jar,  indicating 
the  formation  of  water.  The  burning  of  hydrogen  in  air  is,  in  fact, 
a  chemical  change  in  which  the  oxygen  of  the  air  unites  with  hydro- 
gen to  form  water.  If  care  has  not  been  taken  to  expel  all  the  air 
from  the  Woulf's  bottle  before  collecting  the  hydrogen,  the  mixture 
of  hydrogen  and  air  will  explode  when  a  light  is  brought  to  the 
mouth  of  the  jar. 

The  explosive  character  of  mixtures  of  hydrogen  and  oxygen  may 
be  more  strikingly  demonstrated  by  filling  a  soda-water  bottle  by 
displacement  with  a  mixture  of  one  volume  of  oxygen  and  approxi- 
mately two  volumes  of  hydrogen.  The  bottle  is  wrapped  in  a  cloth 
to  protect  the  hand  in  case  of  accident,  and  on  applying  a  light  to 
the  mouth  a  violent  explosion  occurs.  The  gases  combine  under 
ordinary  conditions  only  when  the  temperature  is  sufficiently  high. 
The  application  of  the  taper  to  the  mouth  of  a  bottle  containing  the 
mixed  gases  leads  to  the  combination  of  a  small  portion  of  the  mixture, 
and  the  heat  given  out  raises  the  rest  of  the  mixture  above  the  tempera- 
ture of  combination,  so  that  the  change  is  practically  instantaneous. 

The  fact  that  water  is  formed  when  hydrogen  burns  in  air  may 
be  illustrated  by  causing  a  jet  of  burning  hydrogen  to  impinge 
against  a  cold  surface,  as  shown  in  Fig.  16.  The  hydrogen  prepared 
in  the  Woulf's  bottle  A  is  dried  by  passing  through  the  U  -tubes 
B  and  C,  which  contain  anhydrous  calcium  chloride,  and  the  burning 
jet  is  in  contact  with  the  flask  D,  which  contains  a  large  quantity 
of  cold  water.  In  a  few  minutes  sufficient  water  will  have  collected 
to  allow  of  its  recognition. 

Hydrogen  and  oxygen  combine  fairly  rapidly  at  room  temperature 
in  the  presence  of  a  catalytic  agent  such  as  platinum  black.  When 
perfectly  dry,  however,  the  mixed  gases  do  not  combine  even  when 
heated  to  the  melting-point  of  silver  (Baker). 


38     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Hydrogen  has  so  great  an  affinity  for  oxygen  that  it  not  onh 
combines  with  the  free  gas,  but  even  removes  it  from  combination 
with  other  elemeHts.  This  may  conveniently  be  shown  by  heating 


FIG.  1 6. 

black  copper  oxide  in  a  bulb  tube  A  (Fig.  17),  and  passing  a  stream 
of  dry  hydrogen  over  it.  In  a  few  minutes  the  black  powder  will  be 
observed  to  turn  red,  and  finally  only  metallic  copper  remains  in 

A  a 


FIG.  17. 

the  bulb.    The  chemical  change  in  this  case  is  represented  by  the 
equation 

Copper  oxide + hydrogen  =  copper  -f  water. 

The  water  may  be  collected  in  the  tube  B  placed  behind  the  bulb 
tube. 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     35 


Oxidation  and  Reduction — It  has  already  been  pointed  out 
that  the  addilion  of  oxygen  to  another  element  or  to  a  chemical  com- 
pound is  termed  oxidation,  and  the  same  term  is  applied  when  hydro- 
gen is  removed  from  a  chemical  compound.  The  removal  of  oxygen 
from  copper  oxide  by  hydrogen,  just  described,  is  called  reduction^  and 
the  same  term  is  applied  to  the  adding  on  of  hydrogen  to  another  ele- 
ment or  compound.  Reduction  is  therefore  the  converse  of  oxidation. 

Occlusion  of  Hydrogen  — Certain  metals,  more  particularly 
platinum,  palladium,  and  iron,  possess  the  remarkable  property  of 
absorbing  or  occluding  many  times  their  own  volume  of  hydrogen 
when  heated,  and  retaining  it  at  the  ordinary  temperature.  The 
following  table  gives  the  maximum  volume  of  hydrogen,  referred 
to  o°  and  760  mm.,  which  can  be  retained  by  one  volume  of  the  finely 
divided  metals  : — 


Palladium  black 
Platinum  sponge 
Gold  precipitated 
Iron  reduced 


873  vols. 

493  ». 
46  „ 
19.2  „ 


Nickel  reduced 
Cobalt       „ 
Copper      „ 
Silver  powder 


18    vols. 

153      » 
4.8    „ 

0-95  » 


The  occluded  hydrogen  does  not  appear  to  enter  into  chemical 
combination  with  the  metals,  but 
the  phenomenon  is  by  no  means 
well  understood.  No  other  gas  is 
occluded  by  metals  in  general  to 
the  same  extent  as  hydrogen. 

Collection  of  Gases  by  Dis- 
placement of  Air— A  very  light 
gas,  such  as  hydrogen,  may  be 
collected  by  downward  displace- 
ment of  air.  The  delivery  tube  is 
passed  nearly  to  the  bottom  of  the 
inverted  jar  (Fig.  18  <z),  and  after 
some  time  it  will  be  found  that 
the  jar  is  full  of  hydrogen.  If 
the  jar  is  kept  inverted,  the  hydro- 
gen will  be  retained  for  some 
time,  but  when  it  is  placed  mouth 
upwards  the  gas  soon  escapes. 
Similarly,  a  gas  heavier  than  air 

(e.g.  chlorine)    can  be   collected   by  upward   displacement   of   air 
(Fig.  18  b\ 


FIG.  18. 


40     A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

GENERAL  PROPERTIES  OF  GASES 

OVeneral — It  is  well  known  that  matter  can  exist  in  three  forms 
or  states  of  aggregation,  the  gaseous,  liquid,  and  solid  states. 
Further,  the  particular  form  in  which  a  definite  substance  occurs 
depends  on  the  external  conditions  ;  thus  when  the  temperature  is 
raised  solid  water  or  ice  changes  to  liquid  water,  and  at  a  still  higher 
temperature  water  changes  to  steam,  which  is  a  gas. 

Solids  have  definite  form,  and  the  volume  of  a  solid  does  not 
alter  greatly  when  the  external  conditions,  such  as  temperature  and 
pressure,  are  altered.  Liquids  differ  from  solids  in  that  the/  readily 
take  the  shape  of  the  vessel  containing  them ;  like  solids  they  have 
a  definite  volume,  which  is  not  greatly  affected  by  changes  of 
temperature  and  pressure.  Gases,  on  the  other  hand,  have  no 
definite  volume;  they  are  characterized  by  their  tendency  to  fill 
completely,  and  to  a  uniform  density,  any  available  space.  This 
fact  cannot  be  illustrated  very  conveniently  with  colourless,  odour- 
less gases,  such  as  hydrogen  and  oxygen,  but  may  readily  be  shown 
with  bromine,  which  is  red  in  colour.  If  a  few  drops  of  liquid 
bromine  are  placed,  by  means  of  a  pipette,  at  the  bottom  of  a  tall 
gas  jar  containing  air,  and  the  jar  is  covered  and  set  aside,  it  will 
be  found  in  course  of  time  that  the  gaseous  bromine,  which  at  first 
was  only  to  be  found  at  the  bottom  of  the  jar,  has  become  uniformly 
distributed  throughout  the  confined  space,  as  shown  by  the  colour. 

In  giving  the  weights  of  a  litre  of  hydrogen  and  of  oxygen  respec- 
tively, the  temperature  and  pressure  under  which  the  gases  were 
measured  have  been  given.  This  implies  that  the  volume  of  a 
definite  quantity  of  a  gas  depends  on  the  conditions  under  which  it  is 
measured,  and  such  is  the  case.  The  laws  expressing  the  behaviour 
of  gases  under  varying  conditions  will  now  be  briefly  considered. 
The  most  remarkable  fact  about  these  laws  is  that  they  are  to  a 
great  extent  independent  of  the  nature  of  the  gas  ;  the  volume  of  all 
gases  is  affected  by  changes  of  temperature  and  pressure  to  much 
the  same  extent. 

Relation  between  Volume  and  Pressure  for  Gases— 
When  the  pressure  upon  a  confined  volume  of  gas  is  increased,  the 
volume  diminishes.  The  exact  relationship  between  pressure  and 
volume  was  discovered  by  Boyle  (i6£i),  and  is  known  as  Boyle's 
law ;  the  law  may  be  formulated  as  follows  :  At  constant  tempera- 
ture the  volume  of  a  given  mass  of  a  gas  is  inversely  proportional  to 
the  pressure  to  which  it  is  subjected.  This  means  that  if  the  pressure 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     41 

on  a  gas  is  doubled,  the  volume  is  halved,  if  the  pressure  is  made 
four  times  as  great,  the  volume  is  reduced  to  one-fourth,  and  so  on. 
Boyle's  law  may  therefore  be  put  in  the  alternative  form,  that  the 
product  of  the  pressure/  and  volume  v  of  a  definite  mass  of  a  gas  is 
constant  at  constant  temperature.  If  v  is  the  volume  at  pressure  p 
and  z/j  the  volume  at  pressure^  then/  v  =/x  v^  —  constant. 

The  validity  of  Boyle's  law  may  be  tested  by  means  of  the  arrange- 
ment represented  in  Fig.  19.     A  straight  tube,  which  must  be  more 


_a 


ABC 

FIG.  19. 

than  80  cm.  long,  and  closed  at  one  end,  is  filled  with  mercury  and 
inverted  in  a  mercury  trough  in  such  a  way  that  no  air  enters.  It  will 
be  found  that  the  mercury  only  fills  part  of  the  tube,  the  upper  un- 
shaded part  is  a  vacuum.  The  mercury  is  supported  in  the  tube  by  the 
pressure  of  the  atmosphere,  and  the  difference  of  level  a  b  between  the 
surface  of  the  mercury  in  the  trough  and  that  in  the  tube  measures 
the  pressure  of  the  atmosphere  (Fig.  19  A).  The  average  height  of  the 
column  of  mercury  is  about  760  mm.  (30  inches).  The  apparatus  is 


42      A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

called  a  barometer.  Another  tube  has  a  long  limb  which  is  open  tc 
the  air  and  a  short  limb  which  is  closed.  A  little  mercury  is  poured 
into  the  open  end,  and  so  adjusted  that  a  column  of  air  is  confined  in 
the  shorter  limb,  while  the  mercury  stands  at  the  same  level  in  both 
limbs  (Fig.  19  B).  Under  these  circumstances  the  confined  air  is 
necessarily  under  atmospheric  pressure,  since  the  atmosphere  is 
pressing  upon  the  surface  of  the  mercury  in  the  open  tube,  and  the 
enclosed  air  upon  that  in  the  closed  tube  ;  the  fact  that  the  mercury 
is  standing  at  the  same  level  in  both  tubes  indicates  the  equality  of 
the  two  pressures.  The  length  of  the  confined  column  of  air  is  then 
measured,  and  mercury  poured  into  the  open  end  of  the  tube  till  the 
difference  in  level  between  the  mercury  surfaces  in  the  two  limbs 
corresponds  with  one  atmosphere  pressure,  as  read  off  on  the 
barometer  (Fig.  19  C).  The  enclosed  air  is  now  under  two  atmo- 
spheres pressure,  and  it  will  be  found  by  measurement  that  its 
volume  has  been  reduced  to  one-half,  in  accordance  with  Boyle's  law. 
If  a  longer  tube  is  taken,  the  validity  of  the  law  can  be  tested  by 
applying  still  higher  pressures.  The  apparatus  with  which  Boyle 
established  the  law  was  very  similar  to  that  just  described. 

Boyle's  law  is  not  strictly  true  for  any  actual  gas,  but  is  very  nearly 
so  for  gases  such  as  hydrogen,  oxygen,  and  nitrogen,  which  are  very 
difficult  to  liquefy.  Under  ordinary  conditions  hydrogen  is  less  com- 
pressible, and  all  the  other  gases  more  compressible,  than  the  law 
indicates. 

Relation  between  Volume  and  Temperature  for  Gases 
— When  a  gas  is  heated  at  constant  pressure  it  expands.  Careful 
experiment  has  revealed  the  remarkable  fact  that  the  increase  in  the 
volume  of  a  gas  for  a  given  rise  of  temperature  is  a  constant, 
independent  of  the  nature  of  the  gas.  If  a  definite  volume  of  a  gas 
is  raised  i°  in  temperature  at  constant  pressure,  it  expands  by  1/273 
of  its  volume  at  o°  ;  if  it  is  raised  10°  in  temperature  it  expands  by 
10/273  of  its  volume  at  o°.  In  the  same  way,  if  a  definite  volume  of 
a  gas  is  cooled  i°,  the  volume  diminishes  by  1/273  of  its  value  at 
o°.  These  statements  are  summarized  in  Charles's  law,  which  may 
be  expressed  as  follows  :  At  constant  pressure  a  gas  expands  or  con- 
tracts by  1/273  of  its  volume  at  o°  for  every  change  in  temperature  oj 
i°  C. 

This  law  may  be  put  in  a  more  concise  form  on  the  basis  of  the 
following  considerations.  Imagine  a  gas  confined  in  a  tube  graduated 
in  c.c.  (Fig.  20)  by  the  air-tight  weightless  piston  (shaded  in  the 
digram)  which  moves  without  friction  in  the  tube,  and  that  the 


240  CC 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     43 

volume  of  the  gas  occupies  273  graduations  of  the  tube  at  o°  C.  If 
now  the  temperature  of  the  gas  is  raised  i°,  the  volume  increases  by 
1/273  of  that  at  o°,  that  is,  by  one  graduation,  and  the  new  volume  is 
274  c.c.  Similarly  if  the  temperature  is  raised  to 

100°,  the  new  volume  is  273  ( i  +  —  J  =  373  c.c.,  and 
if  the  temperature  is  lowered  to-ioo°,  the  new 

volume  is  273  (  i-^))  =  i73  c.c.     It  is  evident, 
\       273/ 

therefore,  that  the  arrangement  in  question  might 
be  used  as  a  thermometer  for  measuring  the  tem- 
perature, as   every    change   of  temperature   of  x 
degrees  corresponds  with  a  change  of  .r  graduations 
on  the  scale.     Further,  if  the  same  rule  continues 
to  hold,  the  vohime  of  the  gas  will  theoretically  be 
zero  when  the  temperature  has  fallen  to  —  273°  C. 
and  no   lower   temperature   can   be   registered  by    200CC 
our  gas  thermometer.     These  considerations  have 
led  to  the  establishment  of  a  new  scale  of  tempera-   /aocc 
ture,  the  so-called  absolute  scale,  the  zero  on  the 
absolute  scale  being -273°  C.     The  temperatures   /6<7CC 
on  the  absolute  scale,  the  so-called  absolute  tem- 
peratures, are  shown  on  the  right  hand  of  the  scale,  H 
and  it  is  evident  that   absolute  temperatures  are    .^^ 
obtained  by  adding  273  to  the  corresponding  tem- 
peratures  on   the    Centigrade    scale.      It  is  now    /ooce 
clear  that  our  constant  pressure  gas  thermometer 
measures   absolute   temperatures,    or   to   put  the 
matter  in  another  way  :  At  constant  pressure,  the 
volume   of  a  gas  is  proportional  to   the  absolute 
temperature.      This  is  the  form  of  Charles's  law 
which  is  most  useful  in  calculations  dealing  with 
gaseous  volumes.      The  method  of  employing  it  will 
now  be  illustrated  by.  two  examples. 

(i)  If  the  temperature  of  a  quantity  of  gas  which 
measures  zoo  c.c.  at  10°  is  changed  to  130°  at  con- 
stant pressure,  what  is  the  new  volume  ? 

As  the  volume  is  proportional  to  the  absolute  temperature,  273  c.c. 
at  o°  C.  will  measure  283  c.c.  at  10°  and  273+130=403  c.c.  at  130°. 

403 
Hence  icoc.c.  at  10°  will  measure  100  x  -§-  =  142.4  c.c.  at  130°. 


60  CC 


^oec 
/occ 
occ' 


S7°C   330° A 


37°C  310  "A 


70  c   zeo°a 

0°  C     273°  A 

-/3°C    260°A 

-23°C  2SO°A 

•  -33°  C   240°A 

--53°C    220°A 
-79°C  200°A 


-113° C  I60°A 


-AJ3°C   MO'A 


-I73°C  /00CA 
-  -I03°C  80°A 
-2/3°C  60°A 
~233"C  40°A 

-2S3°C  tO°A 
-263°C  10° A 
-Z73°C  0°A 

FIG.  20. 


44     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


(2)  What  will  be  the  volume  of  the  same  quantity  of  gas  if  measured 
at -80°  C.?  2830.0.  at  io°C.  will  measure  (2 73 -80)  =193  c.c.  at -80°. 

IQ3 

Hence  the  volume  at -80°  must  be  ioox-~  =68.2  c.c. 

283 

The  important  point  to  be  observed  in  making  the  above  correc- 
tions is  the  proper  placing  of  numerator  and  denominator.  If  the 
volume  at  a  lower  temperature  is  required,  the  smaller  absolute 
temperature  is  of  course  used  as  numerator,  as  the  final  volume 
is  smaller  than  the  initial  volume  ;  if  the  volume  at  a  higher  tem- 
perature is  required,  the  greater  absolute  temperature  is  used  as 
numerator. 

The  expansion  which  unit  volume  of  a  gas  shows  when  raised  i°  in 
temperature  at  constant  pressure  may  be  termed  its  coefficient  of  ex- 
pansion, and  it  has  been  pointed  out  in  this  section  that  the  coefficient 
of  expansion  of  all  gases  is  approximately  the  same,  and  is  equal  to 
1/273  or  0.003665  of  its  volume  at  o°.  The  coefficient  of  expansion  is 
often  indicated  by  the  letter  #,  and  if  v  t  is  the  volume  of  the  gas  at  /° 
C.  and  VQ  its  volume  at  o°  C.  then  vt  =  v0  (i  +a  t}  =  v§  (1+0.003665  f) 
according  to  the  law  of  Charles.  This  equation  may  be  used  to  correct 
gases  for  changes  of  temperature,  but  is  not  so  convenient  as  the 
absolute  temperature  method  already  described. 

So  far,  it  has  been  assumed  that  the  coefficient  of  expansion  of 
all  gases  at  constant  pressure  is  exactly  the  same.  This,  however, 
is  by  no  means  the  case,  the  law  of  Charles,  like  that  of  Boyle,  being 
only  an  approximate  one.  The  deviations  from  the  simple  law  depend 
both  on  the  nature  of  the  gas  and  on  the  temperature  at  which  the 
observations  are  made,  and  for  all  gases  (except  hydrogen)  the  co- 
efficient approaches  the  value  0.003665  the  more  nearly  the  lower  the 
pressure,  and  the  further  the  gas  is  removed  from  its  temperature  of 
liquefaction.  This  is  shown  by  the  following  table — 

Coefficients  of  Expansion  of  Gases  at  Constant  Pressure 


Gas. 

Pressure  at  o° 
in  cm.  Mercury. 

Coefficient  of 
Expansion. 

Gas. 

Pressure  at  o°  in 
cm.  Mercury. 

Coefficient  of 
Expansion. 

Hydro- 

76 cm. 

0.0036613 

Carbon 

76  cm. 

0.003710 

gen 

254    .. 

0.0036616 

diox- 
ide. 

252    „ 

0.003845 

Air 

76    „ 

0.003-  7i 

•  • 

257     ,. 

0.003695 

HYDROGEN— GENERAL  PROPERTIES  OF  GASES     45 

General  Equation  for  Gases— So  far,  we  have  investigated 
the  relationship  between  the  volume  and  pressure  of  a  gas  at  con- 
stant temperature  (Boyle's  law),  and  between  the  volume  and 
temperature  of  a  gas  at  constant  pressure  (Charles's  law).  It  still 
remains  to  discuss  the  relationship  between  pressure  and  tempera- 
ture at  constant  volume. 

A  little  consideration  shows,  however,  that  the  law  in  question 
can  readily  be  deduced  when  Boyle's  law  and  Charles's  law  have 
been  established.  Suppose,  for  example,  we  have  a  definite  quantity 
of  a  gas  contained  in  a  vessel  of  fixed  capacity  at  273°  abs.,  and  raise 
its  temperature  to  546°  abs, ;  if  free  to  expand,  its  volume  would  be 
doubled  at  constant  pressure.  As,  however,  its  volume  is  kept  con- 
stant, its  pressure,  according  to  Boyle's  law,  must  be  doubled.  From 
these  considerations  we  deduce  the  third  of  the  simple  gas  laws  : 

At  constant  volume^  the  pressure  of  a  gas  is  proportional  to  the 
absolute  temperature.  Like  the  other  laws  the  above  law  is  only 
approximately  true. 

It  can  readily  be  shown  that  the  three  gas  laws  can  be  expressed  in 
the  simple  formula 

'constant, 

where /0  and  T/O  are  the  pressure  and  volume  of  a  definite  quantity  of 
gas  at  the  absolute  temperature  7"0,  and  pv  and  ^  the  pressure  and 
volume  at  the  absolute  temperature  7\. 

The  use  of  this  formula  in  finding  the  volume  of  a  gas  when  both 
temperature  and  pressure  vary  may  be  illustrated  by  the  following 
example. 

A  definite  quantity  of  a  gas  measures  500  c.c.  at  20°  and  500  mm. 
pressure,  what  is  its  volume  at  -  30°  and  900  mm.  pressure  ? 

Substituting  in  the  general  equation,  we  have 

500  x  5 oo  _  900  xz/i 
273  +  20      273  -  30 

Whence  ^-K^ISSE*  2-43= 230.4  c.c. 

900X293 

The  same  result  may  also  be  obtained  by  correcting  first  for  the 
change  of  pressure  by  Boyle's  law,  the  new  volume  being  then 
corrected  tor  change  of  temperature  by  Charles's  law.  The  general 


46     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


formula  should  not  be  used  until  the  student  is  also  familiar  with  the 
latter  method  of  calculation. 

The  general  formula,  like  the  laws  of  which  it  is  a  summary,  is 
only  approximately  valid,  but  is  the  more  nearly  true  the  higher 
the  temperature,  the  lower  the  pressure,  and  the  further  the  gas 
is  removed  from  its  temperature  of  liquefaction.  The  theoretical 
bearing  of  this  statement  is  discussed  later. 

Diffusion  of  Gases— It  has  already  been  stated  that  when 
two  gases  are  brought  together  they  ultimately  become  uniformly 
mixed,  even  against  the  force  of  gravity.  This  may  be  shown  by  a 
modification  of  the  experiment  described  on  p.  40.  If  a  few  drops 
of  bromine  are  carefully  placed,  by  means  of  a 
pipette,  at  the  bottom  of  a  jar  of  hydrogen,  and 
the  jar  is  covered,  it  will  be  found  after  a  short 
time  that  the  colour  is  equal  throughout  the  space, 
showing  that  the  bromine  gas  or  vapour,  although 
eighty  times  heavier  than  hydrogen,  has  become 
uniformly  distributed  through  it.  The  same  fact 
can  be  illustrated  by  placing  a  covered  jar  of 
hydrogen  mouth  to  mouth  over  a  covered  jar  of 
the  heavy  greenish-yellow. gas  chlorine,  and  care- 
fully withdrawing  the  covers.1  After  a  time  it  will 
be  noticed  that  the  chlorine  has  become  distributed 
through  the  hydrogen  in  the  upper  jar,  and  on 
bringing  a  light  to  each  jar  an  explosion  will  occur, 
showing  that  both  contain  the  two  gases  in  con- 
siderable proportion.  This  process  of  mixing  is 
termed  the  diffusion  of  gases.  The  same  pheno- 
menon is  observed  when  two  gases  are  separated 
by  means  of  a  porous  plate  ;  in  this  case  it  may  be  assumed  that  the 
gas  particles  readily  pass  through  the  pores,  so  that  diffusion  is  not 
hindered. 

Experiment  shows  that  gases  diffuse  at  very  different  rates,  and 
that  a  gas  diffuses  the  more  rapidly  the  lower  its  density.  This 
may  be  shown  very  instructively  by  means  of  the  apparatus  re- 
presented in  Fig.  21.  A  long  glass  tube,  the  upper  end  of  which 
is  fixed,  by  means  of  a  cork,  into  an  inverted  porous  pot  A  (a 
battery  jar),  is  fitted  with  a  cork  into  one  of  the  openings 
of  a  Woulf's  bottle  C,  which  is  partly  filled  with  water.  Into 

1  As  the  gases  combine  explosively  in  sunlight,  this  experiment  must  be  per- 
formed  in  diffused  daylight. 


FIG.  21. 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     47 

the  other  opening  of  the  bottle  is  fitted  a  cork  carrying  a  short 
glass  tube,  the  lower  end  of  which  dips  in  the  liquid  in  the  bottle, 
the  upper  end  being  drawn  out  to  a  point.  If  now  a  beaker, 
B,  is  placed  over  the  porous  pot  and  hydrogen  is  passed  up  into 
B,  diffusion  of  the  hydrogen  into  the  pot  will  take  place  much 
more  rapidly  than  the  air  diffuses  outwards,  the  pressure  inside 
the  pot  is  considerably  increased,  and  water  is  forced  out  at  the 
narrow  end  of  the  short  tube  in  the  form  of  a  jet.  If  now  the 
beaker  with  the  hydrogen  is  removed,  the  jar  is  again  surrounded 
with  air,  the  mixture  of  gases  in  the  pot  (mainly  hydrogen)  diffuses 
out  more  rapidly  than  the  air  enters,  the  pressure  inside  the  pot 
diminishes,  the  water  falls  in  the  long  tube,  and  bubbles  of  air  are 
drawn  into  the  bottle. 

The  Law  of  Gaseous  Diffusion  —  The  law  expressing  the 
relative  rates  at  which  gases  diffuse  was  first  established  by  Thomas 
Graham,  and  may  be  formulated  as  follows  :  The  relative  rates  of 
diffusion  of  two  gases  are  inversely  proportional  to  the  square  roots 
of  their  densities.  If  the  velocity  of  diffusion  is  represented  by  v,  and 
the  density  of  the  gas  by  d,  Graham's  law  can  be  expressed  by  the 
formula 

v=  J\\d. 

In  order  to  illustrate  the  law,  we  may  consider  the  relative  rates 
at  which  hydrogen  (density  i)  and  chlorine  (density  35.5)  diffuse 
through  a  porous  partition  into  air.  The  ratio  of  the  square  roots 
of  the  densities  is  \/i:V35-5j  approximately  1:6,  hence  the  rela- 
tive rates  of  diffusion,  being  inversely  as  the  square  roots  of  the 
densities,  are  as  6:1. 

The  results  obtained  by  Graham,  and  given  in  the  following  table, 
indicate  that  the  experimental  results  are  represented  very  satis- 
factorily indeed  by  the  above  formula. 


Gas. 

Density,  d,  compared 
with  Air  as  Unity. 

ViK 

Observed  Rate  of 
Diffusion  compared 
with  Air  as  Unity. 

Hydrogen     .     . 
Nitrogen  . 
Oxygen    .     .     . 
Carbon  dioxide 
Sulphur  dioxide 

0.0695 
0.9713 
1.  1056 
1.5290 
2.247 

.3-794 
1.014 
0.951 
0.809 
0.667 

3-83 
1.014 
0.949 
0.812 
0.68 

48     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Not  only  may  the  rate  of  diffusion  of  a  gas  be  calculated  from 
its  density,  but  the  density  may  be  determined  indirectly  by  ob- 
serving the  rate  of  diffusion.  This  principle  has  been  found  useful 
on  more  than  one  occasion  in  chemical  investigations. 

The  differences  in  the  rates  of  diffusion  of  gases  have  been  taken 
advantage  of  to  effect  a  partial  separation  of  the  constituents  of  a 
gaseous  mixture — a  process  known  as  atmolysis.  If,  for  instance, 
a  mixture  of  hydrogen  and  oxygen  is  passed  through  a  long  tube 
of  porous  material,  such  as  a  series  of  tobacco  pipe  stems,  and 
the  issuing  gas  is  collected  over  water,  it  will  be  found  to  be 
much  richer  in  oxygen  than  the  original  mixture. 

The  rate  of  passage  of  gases  through  a  very  small  hole  in  a  plate 
(preferably  a  platinum  plate)  was  also  investigated  by  Graham,  and 
was  found  to  follow  the  same  law  as  gaseous  diffusion.  The  pheno- 
menon is  termed  gaseous  effusion. 

The  Kinetic  Theory  of  Gases— It  is  natural  to  consider 
whether  any  mental  picture  of  the  nature  of  gases  can  be  sug- 
gested which  may  serve  to  account  for  the  simple  laws  which 
have  been  found  to  represent  their  behaviour,  and  also  to 
account  for  the  deviations  from  these  laws.  Such  a  mechanical 
representation  was  brought  forward  by  Bernoulli  as  far  back  as 
1738,  and  has  been  developed  by  later  workers  into  the  kinetic 
theory  of  gases.  According  to  the  theory,  gases  are  made  up  of 
small,  perfectly  elastic  particles  (the  chemical  molecules,  p.  106), 
which  are  in  continual  rapid  motion,  colliding  with  each  other 
and  with  the  walls  of  the  containing  vessel.  The  particles  of  any 
one  gas  are  supposed  to  be  identical,  but  differ  from  the  particles 
of  other  gases  in  respect  to  mass,  speed,  and  other  properties. 
The  space  actually  filled  by  the  gas  particles  is  supposed  to  be 
smaller  than  that  which  they  inhabit  under  ordinary  conditions ; 
they  have,  therefore,  a  comparatively  large  free  space  in  which  to 
move,  and  are  practically  free  from  each  other's  influence  except 
during  a  collision.  The  average  distance  over  which  a  particle 
moves  before  colliding  with  another  particle  is  termed  the  mean 
free  path  of  the  particle. 

According  to  this  theory,  the  pressure  exerted  by  a  gas  on  the 
walls  of  the  containing  vessel  is  due  to  bombardment  by  the 
moving  particles.  It  is  evident  therefore  that  the  magnitude  of 
the  pressure  must  depend  on  the  mass  and  the  velocity  of  the 
particles.  It  can  be  shown  that  the  pressure  exerted  by  a  single 
particle  is  proportional  to  its  mass  and  to  the  square  of  its  velocity, 


HYDROGEN— GENERAL  PROPERTIES  OF  GASES     49 

and  the  total  pressure  is  the  sum  of  the  pressures  exerted  by  each 
particle.  It  should,  however,  be  remembered  that,  owing  to  collisions 
and  for  other  reasons,  the  speed  of  the  particles  in  any  one  gas  is 
probably  by  no  means  uniform,  but  varies  considerably  about  a  mean 
value. 

It  will  now  be  shown  that  the  simple  gas  laws  are  in  full  accord 
with  the  view  as  to  the  constitution  of  gases  just  stated.  If  at 
constant  temperature  the  volume  in  which  a  definite  mass  of  gas  is 
confined  is  halved,  the  number  of  impacts  on  the  walls  in  a  given 
time  is  doubled ;  in  other  words,  the  pressure  of  the  gas  is  doubled. 
This  is  Boyle's  law,  that  the  product  of  the  pressure  and  volume  of  a 
given  mass  of  gas  is  constant  at  constant  temperature. 

Further,  we  have  seen  that  at  constant  volume  the  pressure  of  a 
given  mass  of  gas  is  proportional  to  the  absolute  temperature.  As 
increase  of  temperature  cannot  alter  the  number  of  the  particles,  the 
observed  increase  of  pressure  must,  according  to  the  kinetic  theory, 
be  due  to  an  increase  in  the  speed  of  the  particles,  resulting  in  a 
greater  number  of  impacts  on  the  walls  of  the  vessel  in  a  given  time. 
We  have  already  learnt  that  the  pressure  of  a  gas  is  proportional  to 
the  square  of  the  rectilineal  velocity  of  the  particles,  hence  it  follows 
that  the  square  of  the  rectilineal  velocity  of  the  particles  is  proportional 
to  the  absolute  temperature. 

These  considerations  throw  an  interesting  light  on  the  physical 
meaning  of  the  absolute  zero.  As  the  temperature  falls,  the  rectilineal 
velocity  of  the  particles  steadily  diminishes  and  finally,  at  the  abso- 
lute zero,  they  must  theoretically  come  to  rest.  The  absolute  zero  is 
therefore  the  lowest  temperature  theoretically  attainable.  In  actual 
practice  it  has  never  been  reached,  but  in  his  recent  investigations  on 
the  liquefaction  of  helium  (p.  210),  Kammerlingh  Onnes  has  got  within 
2°  of  the  absolute  zero.  On  the  other  hand,  there  is  no  theoretical 
upper  limit  to  the  speed  of  the  particles,  and  therefore  no  upper  limit 
of  temperature. 

Not  only  does  the  kinetic  theory  of  gases  afford  a  satisfactory 
interpretation  of  the  gas  laws,  but  it  also  accounts  for  the  more 
important  deviations  from  these  laws.  In  deducing  Boyle's  law, 
we  have  tacitly  assumed  that  when  the  volume  is  halved  the  space 
in  which  the  particles  move  has  been  halved  ;  in  other  words,  we 
have  neglected  the  space  filled  by  the  particles  in  comparison  with 
that  which  they  inhabit.  Further,  we  have  made  no  allowance  for  a 
possible  attraction  between  the  particles.  If  such  an  attraction 
exists,  it  must  be  the  greater  the  nearer  the  particles  approach  each 
4 


50     A  TEXT-BOOK   OF    INORGANIC   CHEMISTRY 

other.  Hence,  when  the  pressure  is  doubled,  the  effect  of  the 
attraction,  which  must  lead  to  a  diminution  of  volume,  is  super- 
imposed on  the  regular  contraction  according  to  Boyle's  law,  so  that 
the  value  of  v,  and  hence  the  product,  pv,  is  less  than  the  calculated 
value. 

On  the  other  hand,  a  little  consideration  shows  that  the  effect  of  the 
finite  size  of  the  particles  is  such  that  the  value  of  pv  tends  to  increase 
with  increasing  pressure.  Let  us  assume  that  a  particle  is  moving 
backwards  and  forwards  between  the  walls  a  and  b  and  that  the 
distance  between  the  walls  is  20  times 


the  diameter  of  the  particle.     The  dis-       | 

tance  traversed  between  each  impact   is   clearly    19  diameters.     If 

now  the  distance  between  the  walls  is  halved,  the  distance  traversed 

between  each  impact  is  only  9  diameters,  and  therefore  the  number 

of  impacts — and  the  pressure — is  more  than  doubled  by  halving  the 

volume. 

It  is  evident  from  the  foregoing  that  the  two  causes  which  bring 
about  the  deviations  from  the  gas  laws  act  in  opposite  directions. 
For  most  gases,  the  effect  due  to  the  attraction  of  the  particles  is 
greater  at  ordinary  pressures  than  that  due  to  the  finite  size  of  the 
particles,  and  therefore  pv  diminishes  at  first  as  the  pressure  is 
increased  (p.  42),  but  at  high  pressures  increases  with  the  pressure. 
For  hydrogen,  however,  the  effect  of  the  volume  correction  counter- 
balances from  the  first  the  attraction  correction,  and  pv  increases 
continuously  with  the  pressure. 

The  kinetic  theory  also  accounts  satisfactorily  for  the  fact  that 
the  deviations  from  the  gas  laws  are  the  smaller  the  lower  the 
pressure.  Under  these  circumstances,  the  attraction  between  the 
particles  becomes  negligible  owing  to  their  distance  apart,  and 
the  volume  of  the  particles  is  negligible  in  comparison  with  the 
total  volume. 

As  already  mentioned,  no  actual  gas  behaves  exactly  according  to 
the  simple  gas  laws,  but  the  deviations  in  the  case  of  hydrogen, 
helium,  nitrogen,  and  other  gases  which  are  difficult  to  liquefy  are 
very  slight  at  ordinary  temperatures  and  pressures..  A  gas  which 
would  follow  the  gas  laws  accurately  is  termed  a  perfect  or  ideal  gas. 

In  the  foregoing  it  has  been  implicitly  assumed  that  the  particles 
themselves  are  incompressible  ;  the  diminution  in  volume  on  com- 
pressing a  gas  is  ascribed  entirely  to  a  diminution  in  the  free  space 
between  the  particles. 


CHAPTER  VI 
WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS 

TTistory — Water  was  for  a  long  time  regarded  as  an  element. 
-*•  -*•  The  discovery  that  it  is  in  fact  a  chemical  compound  of 
"  inflammable  air"  (hydrogen)  and  of  "  dephlogisticated  air"  (oxygen) 
was  made  by  Cavendish  in  1781,  and  he  showed  at  the  same  time 
that  the  gases  combine  in  the  proportion  of  two  volumes  of  hydrogen 
to  one  volume  of  oxygen.  The  method  employed  by  Cavendish  was 
briefly  as  follows.  A  large  graduated  vessel  was  filled  with  a  gaseous 
mixture  containing  approximately  one  volume  of  oxygen  to  two 
volumes  of  hydrogen,  and  was  connected  by  a  bent  tube  to  a  glass 
globe  provided  with  a  brass  stopcock  and  with  two  sealed-in  platinum 
wires,  between  which  an  electric  spark  could  be  passed.  The  glass 
globe,  which  had  previously  been  exhausted  by  the  air-pump,  was 
filled  with  the  gaseous  mixture,  and  the  latter  then  exploded  by  a 
spark.  It  was  then  noticed  that  the  walls  of  the  previously  dry  globe 
were  covered  with  moisture,  showing  the  formation  of  water.  The 
globe  was  again  filled  from  the  graduated  vessel  and  the  mix- 
ture exploded,  the  process  being  repeated  until  sufficient  water 
was  collected  to  put  its  identity  beyond  doubt.  As  practically 
no  gas  remained  in  the  apparatus  after  the  experiment,  the  evidence 
was  conclusive  that  two  volumes  of  hydrogen  and  one  volume  of 
oxygen  combine  to  form  water.  Cavendish's  discovery  as  to  the 
composition  of  water  was  confirmed  by  Lavoisier  (1783),  who  passed 
steam  through  a  heated  iron  tube  and  collected  and  measured  the 
hydrogen. 

Decomposition  of  'Water  into  its  Elements— The  methods 
by  which  it  m%y  be  proved  that  water  is  a  compound  substance  have 
already  been  considered  when  dealing  with  hydrogen.  It  has  been 
shown  that  hydrogen  is  readily  obtained  from  water  when  the  latter 
is  brought  into  contact  with  a  substance  having  a  greater  attraction 
for  oxygen  than  the  latter  has  for  hydrogen.  On  the  same  principle, 
free  oxygen  may  be  obtained  from  water  by  bringing  it  in  contact 
with  a  substance  which  readily  enters  into  chemical  combination  with 

5* 


52     A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 


hydrogen.  A  suitable  substance  for  the  p.urpose  is  gaseous  chlorine. 
When  steam  and  chlorine  are  passed  through  a  red-hot  tube,  free 
oxygen  and  a  chemical  compound  of  hydrogen  and  chlorine  are 
formed,  and  the  oxygen  may  be  collected  over  water  in  the  usual  way. 
This  important  action  is  further  referred  to  in  connexion  with  chlorine 

(P-  87). 

The  direct  decomposition  of  water  into  its  elements  by  the  action 
of  heat  or  of  electrical  energy  has  also  been  referred  to.  The  amount 
of  decomposition  which  steam  undergoes  on  heating  increases  with 
the  temperature,  and  quite  recently  the  variation  of  the  decomposition 
with  the  temperature  has  been  determined  with  considerable  accuracy. 
Some  of  the  results  are  given  in  the  accompanying  table — 


Temperature      .        . 
Amount  of  decomposi 
tion  per  cent. 


1027 

Q 


1207° 
0.0189 


1288° 
0.034 


i882c 
1.18 


1.77 


The  reaction  is  a  reversible  one,  being  represented  by  the  equation 
water^±hydrogen  +  oxygen, 

and  the  above  results  show  that  increase  of  temperature  favours  the 

reaction  represented  by  the  upper  arrow. 

The  Composition  of  Water  by  Synthesis— The  exact  ratio 
by  volume  in  which  hydrogen  and  oxygen  unite 
together  to  form  water  can  readily  be  determined 
by  a  slight  modification  of  Cavendish's  original 
method.  The  apparatus  for  this  purpose  is  illus- 
trated in  Fig.  22.  It  consists  essentially  of  a 
graduated  tube,  called  a  eudiometer,  provided  at  its 
upper  end  with  two  platinum  wires  sealed  through 
the  glass,  the  ends  inside  the  tube  being  so  place*? 
that  an  electric  spark  may  be  passed  between  them. 
The  tube  is  first  filled  with  mercury,  inverted  in  the 
mercury  trough,  and  a  few  c.cs.  of  pure  oxygen,  ob- 
tained by  heating  potassium  permanganate  (p.  22), 
is  passed  up  into  it.  The  volume  ofthe  oxygen  is 
carefully  read  off  on  the  graduated  scale.  In  order 
that  the  volume  of  the  oxygen  may  be  ascertained 
tinder  conditions  comparable  with  that  of  the  hydro- 
gen, its  temperature  and  pressure  must  be  known  ; 

the  temperature  is  read  off  on  the  thermometer  in  the  neighbourhood 

of  the  eudiometer,  and  the  pressure  is  obtained  by  subtracting  the 


FIG.  22. 


WATER— PHYSICAL    PROPERTIES   OF   LIQUIDS     53 

height  of  the  column  of  mercury  in  the  eudiometer  tube  from  that 
of  the  barometer  at  the  time  of  the  experiment.  Since  the  pressure 
of  the  atmosphere  is  balanced  by  the  pressure  of  the  oxygen  +  that  of 
the  mercury  column  in  the  tube,  it  is  evident  that  the  pressure  of  the 
oxygen  is  obtained  as  just  stated.  The  volume  of  the  oxygen  is  then 
corrected  to  normal  temperature  and  pressure  by  the  method  already 
described. 

A  quantity  of  hydrogen,  the  corrected  volume  of  which  is  four  to 
five  times  that  of  the  oxygen,  is  then  introduced,  the  volume  of  the 
mixed  gases  ascertained,  and  corrected  to  normal  temperature  and 
pressure  as  before. 

The  tube  is  then  firmly  pressed  down  below  the  mercury  in  the 
trough  upon  a  plate  of  caoutchouc,  and  the  gases  ignited  by  a  spark 
from  an  induction  coil.  A  bright  flame  is  seen  to  pass  down  the  tube, 
and  on  cautiously  raising  the  end  of  the  tube  from  the  pad,  a  con- 
siderable amount  of  mercury  will  enter.  After  the  temperature  has 
fallen  to  that  of  the  atmosphere,  the  volume  of  the  residual  gas  (un- 
combined  hydrogen)  is  read  off  and  reduced  to  normal  conditions. 

The  mode  of  calculating  the  combining  volumes  of  oxygen  and 
hydrogen  from  these  data  is  best  illustrated  by  an  example. 

Corrected  volume  of  oxygen        ....  31.82  c.c. 

„              „        „  mixed  gases         .        .        .  185.75  c.c. 

„              „         „  residual  hydrogen       .        .  90.35  c.c. 
Hence 

Total  volume  of  hydrogen  taken  185.75  —  31.82  =  153.93  c.c. 

Volume  of  hydrogen  which  has  )  _  =6     g  cc 

combined  with  oxygen  ) 

Therefore 

Volume  of  Oxygen  :  Volume  of  Hydrogen  as  31.82  : 63.58  or  I  :  1.998. 

In  accurate  experiments  a  number  of  corrections  must  be  applied 
which  have  not  been  referred  to  in  the  above  brief  sketch  of  the 
process.  For  example,  the  volume  of  the  water  formed,  although 
very  small  in  comparison  with  that  of  the  gases,  is  not  entirely 
negligible.  According  to  Morley,  who  employed  all  conceivable 
precautions,  the  true  combining  ratio  is  i  :  2.00269  at  o°. 

Volumetric  Composition  of  Water  by  Analysis —  The 
ratio  in  which  hydrogen  and  oxygen  are  present  in  water  can  also 
be  determined  by  electrolyzing  water  between  platinum  electrodes 


54     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

and  collecting  the  gases  separately  (p.  14).  The  volume  of  the 
oxygen  is  always  rather  less  than  half  that  of  the  hydrogen  for  two 
reasons:  (i)  oxygen  is  more  soluble  in  water  than  hydrogen;  (2) 
a  denser  modification  of  oxygen,  termed  ozone,  is  formed  in  small 
amount  at  the  anode. 

Volumetric  Composition  of  Steam  —  If  the  apparatus  in 
which  the  mixture  of  hydrogen  and  oxygen  is  exploded  be  kept  at 
such  a  temperature  that  the  water  produced  remains  in  the  form  of 


FIG.  23. 

steam,  it  is  found  that  two  volumes  of  hydrogen  and  one  volume  of 
oxygen  give  rise  to  two  volumes  of  steam  when  all  the  gases  are 
measured  under  the  same  conditions.  This  may  readily  be  shown  by 
means  of  the  apparatus  represented  in  Fig.  23.  It  consists  essentially 
of  a  U-tube,  one  limb  of  which  acts  as  a  eudiometer  tube,  the  other  is 
open  to  the  air.  A  mixture  of  hydrogen  and  oxygen  in  the  propor- 
tions to  form  water,  prepared  by  electrolysis,  is  introduced  into  the 
eudiometer  tube,  which  is  surrounded  by  a  wide  tube  through  which 
the  vapour  of  a  high-boiling  liquid  may  be  passed.  When  the  vapour 
from  the  boiling  liquid  in  the  flask  has  been  passed  round  the 
eudiometer  tube  sufficiently  long  to  ensure  that  the  temperature  is 
constant,  the  pressure  is  adjusted  to  that  of  the  atmosphere  by  altering 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     55 

the  quantity  of  mercury  in  the  open  limb,  the  volume  of  the  mixed 
gases  is  carefully  noted,  and  the  mixture  then  exploded  by  an  electric 
spark.  When  the  pressure  is  again  adjusted  to  that  of  the  atmosphere 
by  pouring  mercury  into  the  open  limb  until  the  level  is  the  same  in 
the  two  tubes,  it  will  be  found  that  the  volume  of  the  gas  in  the 
eudiometer  has  been  reduced  by  one-third. 

Amyl  alcohol,  boiling-point  132°,  may  conveniently  be  used  as 
jacketing  vapour ;  it  is  condensed  on  leaving  the  tube,  as  shown  in 
the  figure.  Steam  under  increased  pressure  may  also  be  used  for 
the  same  purpose,  although  less  advantageously. 

Gravimetric  Composition  of  Water  — The  composition  of 
water  by  weight  may  be  determined  (a)  from  the  densities  and  com- 
bining volumes^of  the  gases  (£)  directly. 

(a}  According  to  Morley,  the  relative  densities  of  hydrogen  and 
oxygen  are  as  I  :  15.9  and  the  combining  volumes  2.0027  : 1.  Hence 
the  relative  weights  of  hydrogen  and  oxygen  which  combine  to  form 
water  are — 

Hydrogen  2.00269  x  i=  2.0027. 
Oxygen       i  x  15.90x3  =   15.900. 

17.9027.      , 

Hence  17.9027  parts  of  water  contain  2.0027  parts  of  hydrogen  and 
15.900  parts  of  oxygen  ;  otherwise  expressed,  water  contains  11.186 
per  cent,  of  hydrogen  and  88.814  per  cent,  of  oxygen. 

(b]  Direct  Method—  The  method  which  has  been  most  largely 
employed  for  the  direct  determination  of  the  composition  of  water  by 
weight  depends  on  the  reduction  of  copper  oxide  (p.  38).  A  known 
weight  of  the  oxide  is  heated  and  pure  dry  hydrogen  passed  over  it  ; 
the  water  formed  is  collected  in  suitable  vessels  which  are  weighed 
before  and  after  the  experiment.  The  bulb  containing  the  copper 
oxide  is  also  weighed  after  the  experiment  in  order  to  find  the  weight 
of  oxygen  which  is  contained  in  the  amount  of  water  formed.  The 
weight  of  the  hydrogen  is  determined  by  difference. 

An  elaborate  investigation  of  the  composition  of  water  by  this 
method  was  carried  out  by  Dumas  and  Stas  in  1843.  The  arrange- 
ment used  is  illustrated  in  Fig.  24.  Hydrogen,  generated  from  zinc 
ancfsulphuric  acid  in  the  bottle  A,  and  carefully  purified  and  dried  by 
means  of  the  reagents  in  the  U  -tubes,  is  passed  over  the  thoroughly 
dried  copper  oxide  heated  in  B,  the  water  formed  being  collected  in 
the  vessel  C  and  the  three  succeeding  drying  tubes.  From  the  weight 
of  the  vessel  and  drying  tubes  before  and  after  the  experiment,  the 


56     A  TEXT-BOOK  OF   INORGANIC   CHEMISTRY 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     57 

weight  of  the  water  formed  was  obtained,  and  from  the  loss  in  weight 
of  the  copper  oxide  tube  the  quantity  of  oxygen  which  went  to  its 
formation.  As  a  result  of  their  experiments,  Dumas  and  Stas  found 
that  2  parts  by  weight  of  hydrogen  combines  with  15.96  parts  by  weight 
of  oxygen,  a  result  which  does  not  agree  very  well  with  modern  results. 
In  the  hands  of  W.  A.  Noyes,  however  (1889),  this  method  yielded 
results  very  close  to  the  value  now  generally  accepted.  An  improve- 
ment subsequently  introduced  was  to  weigh  the  hydrogen  as  well  as 
the  oxygen  and  the  water  ;  this  was  effected  by  absorbing  a  consider- 
able weight  of  hydrogen  in  palladium,  heating  to  drive  off  the  gas, 
which  was  then  led  over  the  copper  oxide,  the  tube  containing  the 
palladium  being  weighed  before  and  after  the  experiment. 

Morley  (1895)  weighed  the  hydrogen  occluded  in  palladium,  and 
oxygen  in  the  gaseous  form,  burned  the  hydrogen  in  the  oxygen,  and 
weighed  the  water  formed. 

The  result  obtained  by  Morley,  which  is  practically  the  mean  of  the 
most  trustworthy  results  of  other  observers,  is  that  water  contains 
2  parts  by  weight  of  hydrogen  to  15.879  parts  by  weight  of  oxygen,  or 
11.186  per  cent,  of  hydrogen  to  88.814  per  cent,  of  oxygen,  in  exact 
agreement  with  the  result  obtained  from  the  densities  and  combining 
volumes. 

The  fact  established  in  the  foregoing  sections,  that  pure  water, 
whatever  its  source  or  mode  of  preparation,  has  invariably  the  same 
composition,  is  a  matter  of  the  utmost  importance.  We  shall  see 
later  that  the  same  is  true  of  all  other  definite  chemical  compounds. 
This  important  result  is  known  as  the  law  of  constant  composition, 
and  may  be  stated  as  follows :  A  definite  chemical  compound  always 
contains  the  same  elements  in  the  same  proportions  by  weight. 

Physical  Properties  of  Water— Water  at  ordinary  tempera- 
tures is  an  odourless,  tasteless  liquid,  colourless  in  thin  layers,  but  in 
thick  layers  showing  a  slight  bluish-green  colour.  The  colour  is  most 
readily  seen  by  looking  at  a  white  object  through  a  column  of  pure 
water  several  yards  in  length,  contained  in  a  tube  with  blackened 
sides.  The  striking  blue  colour  of  certain  Swiss  lakes,  fed  by  glacier 
streams,  is  probably  due  to  the  intrinsic  colour  of  the  water.  It 
should,  however,  be  remembered  that  a  very  finely-divided  solid 
suspended  in  water  also  gives  rise  to  a  blue  colour  under  certain 
conditions. 

As  is  well  known,  the  freezing-point  of  water  is  taken  as  the  zero 
point,  o°,  and  the  boiling-point  under  760  mm.  pressure  as  100°  on  the 
Centigrade  scale. 


58     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


The  compressibility  of  water  is  very  small.  20,000  volumes  are 
reduced  to  19,999  volumes  by  increasing  the  pressure  by  one  atmos- 
phere. The  effect  of  change  of  temperature  on  the  volume  of  water 
is  also  comparatively  small,  as  is  shown  in  the  accompanying  table,  in 
which  the  volume  and  relative  density  (specific  gravity)  of  water  at 
temperatures  from  o°  to  20°  are  given,  referred  to  water  at  4°  as  unit. 


Tempera- 
ture. 

Volume. 

Rel.  Density. 

Tempera- 
ture. 

Volume. 

Rel.  Density. 

0 

I.OOOI22 

0.999878 

5 

1.000008 

0.999992 

I 

I.OOOO67 

0-999933 

6 

1.003031 

0.999969 

2 

1.000028 

0.999972 

8 

1.000118 

0.999882 

'3 

1.000007 

0.999993 

10 

1.000261 

0-999739 

4 

I.OOOOOO 

I.OOOOOO 

20 

1.001730 

0.998270 

The  figures  in  the  table  show  the  remarkable  fact  that  water  attains 
its  greatest  density  at  a  temperature  in  the  neighbourhood  of  4°.  In 
other  words,  when  water  is  warmed  from  o°,  it  contracts  till  the 
temperature  reaches  4°  (accurately  3.945°)  and  beyond  that  point 
expands  with  rising  temperature. 

Water  expands  on  freezing,  one  volume  of  water  at  o°  giving  1.09082 
volumes  of  ice  at  the  same  temperature.  This  fact  is  of  great  import- 
ance in  effecting  the  disintegration  of  rocks,  the  water  penetrating  into 
cracks  and  exerting  enormous  pressure  on  solidification.  The  burst- 
ing of  water-pipes  in  winter  is  due  to  the  same  cause. 

Water  is  a  very  bad  conductor  of  heat.  If  a  piece  of  ice  be  held 
in  the  lower  part  of  a  test-tube  by  a  piece  of  wire  gauze,  the  water  in 
the  upper  part  of  the  tube  may  be  boiled  without  melting  the  ice. 

The  three  properties  of  water  just  mentioned  are  of  enormous 
importance  in  Nature.  When  a  large  surface  of  water,  such  as  a 
lake,  is  subjected  to  a  low  temperature  the  upper  layers  are  first 
cooled  ;  they  become  denser  and  sink  to  the  bottom.  A  continuous 
circulation  is  thus  set  up,  the  cooler  layers  sinking  and  the  warmer 
rising,  until  the  whole  mass  has  fallen  to  4°.  On  further  cooling,  the 
upper  layers  become  lighter  and  remain  on  the  top  till  they  solidify. 
As,  however,  ice  is  less  dense  than  water,  it  remains  on  the  surface,  so 
that  only  the  upper  layers  solidify.  The  lower  layers  are  further  pro- 
tected against  cooling  by  the  very  small  conductivity  of  the  upper 
layers.  If  it  were  not  for  the  factors  just  enumerated,  more  particu- 
larly the  existence  of  a  point  of  maximum  density  for  water  above 
zero,  the  mass  of  \?ater  would  solidify  as  a  whole,  the  heat  of  summer 
would  be  quite  insufficient  to  melt  the  accumulations  of  ice  formed  in 


WATER— PHYSICAL  PROPERTIES   OF   LIQUIDS     59 

winter,  and  the  climate  of  a  large  part  of  Europe  would  approach  that 
of  the  Polar  regions. 

The  latent  heat  of  fusion  of  ice  is  about  80  calories l  per  gram,  in 
other  words,  80  calories  must  be  supplied  in  order  to  change  one 
gram  of  ice  at  o°  to  water  at  the  same  temperature. 

The  latent  heat  of  vaporization  of  water  is  537  calories,  that  is,  537 
calories  must  be  supplied  in  order  to  change  one  gram  of  water  at 
1 00°  to  steam  at  the  same  temperature. 

Natural  Waters— Owing  to  the  very  great  solvent  power  of 
water,  it  is  never  found  pure  upon  the  earth,  but  always  contains 
dissolved  substances  in  larger  or  smaller  amount.  According  to  the 
amount  of  dissolved  substances  they  contain,  natural  waters  are 
roughly  classified  into — (i)  Fresh  waters,  in  which  the  proportion  of 
substances  in  solution  is  relatively  small ;  (2)  Mineral  waters,  in 
which  the  dissolved  impurities  are  perceptible  to  the  taste. 

Fresh  waters  are  sometimes  divided,  according  to  their  origin,  into 
rain,  river  and  spring  waters,  and  each  of  these  classes  will  now  be 
briefly  considered. 

Rain  Water  is  the  purest  form  of  natural  water.  The  only 
impurities  it  contains  in  appreciable  amount  when  collected  in  the 
•country  are  ammonium  salts,  nitrates,  and  traces  of  organic  matter, 
which  it  takes  up  from  the  atmosphere.  Rain  water  collected  in 
towns  is  much  less  pure,  owing  to  contamination  of  the  atmosphere ; 
it  often  contains  traces  of  free  sulphuric  acid,  resulting  from  the  slow 
oxidation  of  sulphur  compounds.  The  amount  of  solid  matter  in  rain- 
water collected  in  the  country  is  very  variable,  being  greater  at  the 
beginning  than  at  the  end  of  a  shower  owing  to  the  gradual  removal 
of  the  impurities  from  the  atmosphere.  The  average  amount  of 
dissolved  matter  does  not  exceed  0.03  to  0.04  parts  per  1000  of  water 
(0.03  to  0.04  grams  per  litre). 

Spring  Waters  are  always  less  pure  than  rain  water,  owing  to 
their  solvent  action  on  the  strata  through  which  they  pass.  The 
nature  of  the  dissolved  substances  naturally  depends  upon  the 
chemical  composition  of  the  strata,  but  most  fresh  spring  waters 
contain  the  sulphates,  carbonates,  chlorides  and  silicates  of  mag- 
nesium, calcium,  iron,  potassium  and  sodium  in  very  varying  propor- 
tions, and  also  dissolved  gases,  more  particularly  carbon  dioxide. 
The  proportion  of  salts  in  spring  waters  varies  from  0.05  to  3 
grams  per  litre.  As  already  indicated,  spring  waters  containing  an 

1 A  calorie  is  the  amount  of  heat  required  to  raise  i  gram  of  water  i°  C.  in 
temperature. 


60    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

amount  of  dissolved  substances  perceptible  to  the  taste  are  termed 
mineral  waters. 

River  Waters  usually  contain  less  dissolved  matter  than  spring 
waters  ;  but,  owing  to  the  fact  that  rivers  are  mainly  fed  by  surface 
drainage,  they  contain  more  organic  matter,  suspended  and  dissolved, 
than  spring  waters.  The  mineral  substances  in  river  waters  consist 
largely  of  calcium  salts,  more  particularly  the  carbonate  and  sulphate. 
The  proportion  oT  salts  varies  from  o.oc;  to  1.6  grams  per  litre.  The 
average  amount  of  total  solids  in  grams  per  litre  in  some  well-known 
rivers  is  as  follows :  Neva,  0.055  ;  Dee  at  Aberdeen,  0.057  ;  Thames, 
(upper  part)  0.307,  (lower  part)  1.617  ;  Nile,  1.580. 

The  composition  of  the  water  of  lakes  varies  enormously.  The 
water  of  Loch  Katrine,  from  which  the  city  of  Glasgow  is  supplied 
with  water,  contains  about  0.03  grams  of  solid  matter  per  litre,  that  of 
the  Dead  Sea  240  grams  per  litre. 

Mineral  Waters— Sea  water,  though  not  usually  classed  as  a 
mineral  water,  may  appropriately  be  mentioned  here.  When  collected 
far  from  land,  the  composition  of  sea  water  is  very  constant ;  it  con- 
tains about  36  grams  of  solid  matter  per  litre.  The  chief  salt  present 
is  sodium  chloride,  but  magnesium,  potassium,  and  calcium  salts  are 
also  present,  mainly  as  chlorides,  bromides,  and  sulphates.  In  the 
water  of  the  Irish  Sea  the  proportion  of  the  more  important  salts  in 
grams  per  litre  is  approximately  as  follows  (Thorpe,  1870)  :  sodium, 
as  chloride,  27  grams  ;  magnesium,  as  chloride,  3.2  grams,  as  sulphate, 
2  grams,  as  bromide,  0.07  grams  ;  potassium,  as  chloride,  0.75  grams  ; 
calcium,  as  sulphate,  1.4  grams. 

The  typical  mineral  waters  are  classified  according  to  their  most 
important  constituents.  The  chief  types  are  as  follows :  carbonated 
waters,  which  contain  a  large  proportion  of  dissolved  carbon  dioxide — 
examples,  Seltzer,  Apollinaris ;  (2)  alkaline  waters  contain  much 
sodium  bicarbonate — example,  Vichy  ;  (3)  saline  waters  contain  other 
salts  than  sodium  bicarbonate.  Chalybeate  waters  contain  iron  salts 
in  solution;  sulphur  waters  contain  hydrogen  sulphide  and  alkali 
sulphides,  as  at  H arrogate ;  the  wells  at  Epsom  contain  chiefly 
magnesium  sulphate,  the  well-known  Epsom  salts,  etc.  Certain  of 
the  mineral  waters  are  cold  as  they  escape  from  the  ground,  others, 
as  at  Carlsbad,  are  hot. 

Potable  Waters — Water  used  for  drinking  purposes  should  be 
colourless,  odourless,  and  free  from  materials  injurious  to  health. 
The  salts  usually  found  in  river  and  spring  waters  are  not  detri- 
mental ;  in  fact  the  presence  of  them  in  small  proportion  is  advan- 
tageous. The  contamination  of  drinking  water  most  to  be  feared  is 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     61 

the  presence  of  the  germs  of  various  diseases  such  as  cholera  and 
typhoid  fever,  which  often  reach  water  in  sewage ;  and  the  organic 
substances  present  in  sewage  may  also  yield  dangerous  products. 

It  is  evident  therefore  that  bacteriological  examination  of  water 
used  for  drinking  purposes  is  essential.  A  chemical  examination  is 
also  of  importance,  not  so  much  on  account  of  danger  from  salts  that 
may  be  present,  but  because  the  presence  of  certain  constituents  in 
large  proportion  indicates  comparatively  remote  contamination  with 
sewage.  The  presence  of  appreciable  amounts  of  nitrogen,  com- 
bined in  complex  organic  compounds  (albuminoid  nitrogen),  indicates 
recent  contamination.  In  course  of  time,  however,  this  form  of 
nitrogen  becomes  oxidized  to  nitrates  and  nitrites,  and  the  pre- 
sence of  these  salts,  of  sodium  chloride,  and  of  free  ammonia,  above 
a  certain  small  proportion  in  a  sample  of  water  is  usually  a  sign  of 
previous  contamination. 

It  often  happens  that  river  and  spring  waters  not  entirely  free 
from  pollution  have  to  be  used  for  drinking  purposes.  Such  waters 
may  be  rendered  safe  by  boiling,  which  destroys  bacteria,  but  the 
same  object  is  generally  attained  on  the  large  scale  by  filtration. 
This  process  consists  in  passing  the  water  through  beds  of  sand  or 
gravel  with  free  exposure  to  air,  by  which  means  the  bacteria  are 
retained  and  the  organic  matter  oxidized  to  comparatively  harmless 
substances.  The  filter-beds  must  of  course  be  renewed  from  time 
to  time. 

Ozone  is  now  sometimes  employed  for  the  purification  of  water, 
as  it  rapidly  oxidizes  organic  matter  and  can  afterwards  be  readily 
removed. 

On  the  small  scale,  a  filter  of  porcelain  (unglazed  earthenware),  the 
so-called  Pasteur-Chamberland  Filter,  may  be  used ;  in  this  case  the 
water  has  to  be  forced  through  under  pressure.  Filters  of  powdered 
charcoal  are  also  in  use. 

Distillation  of  "Water— Water  can  be  freed  from  most  of  the 
ordinary  impurities  by  converting  it  into  steam,  which  is  then  passed 
through  a  cooling  arrangement  and  condensed  again  to  water.  This 
process  is  known  as  distillation.  An  apparatus  which  can  be  used 
for  this  purpose  is  shown  in  Fig.  25.  The  water1  is  boilrl  in  the 
flask  A,  and  the  steam  passes  along  the  inner  tube  of  the  "  Liebig's 
condenser"  B,  where  it  is  cooled  by  a  stream  of  cold  water  passing 
through  the  outer  tube  of  the  condenser.  The  condensed  water 
collects  in  the  flask  C,  which  is  termed  the  receiver.  The  first 

1  A  purer  distillate  is  obtained  if  a  little  potassium  permanganate  and  alkali  is 
added  to  the  water  before  distilling. 


62     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

portion  of  the  distillate,  which  contains  most  of  the  dissolved  gases, 
is  rejected,  and  the  last  portion  of  the  water,  which  contains  the 
non-volatile  impurities,  is  left  in  the  distilling  flask,  so  that  the  dis- 
tilled water  is  much  purer  than  the  original  sample.  Steam  has, 
however,  a  slight  solvent  action  on  the  glass  of  the  condenser,  and 
in  order  to  avoid  this  source  of  contamination,  condensers  made  of 
materials  not  acted  on  by  steam,  such  as  tin  or  platinum,  are 
largely  used. 

Water  may  also  be   purified  by   partial  freezing,  the  impurities 


FIG.  25. 

remaining  in  the  fluid  portion.  The  ice  is  separated,  allowed  to 
melt,  and  the  process  repeated  if  necessary  (Nernst). 

The  best  method  for  determining  the  purity  of  a  sample  of  water 
is  to  measure  its  electrical  conductivity  (p.  260). 

Distilled  water  has  a  flat  taste,  owing  to  the  absence  of  the 
dissolved  gases  (chiefly  air  and  carbon  dioxide)  which  impart  to 
ordinary  water  a  refreshing  taste. 

SOME  GENERAL  PROPERTIES  OF  LIQUIDS 

As  water  is  in  many  respects  a  typical  liquid,  it  will  be  con- 
venient to  deal  here  with  certain  general  properties  of  liquids, 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS    63 

illustrated  mainly  by  reference  to  the  properties  of  water.  The 
great  majority  of  pure  substances  which  exist  in  the  liquid  form 
can  also  be  obtained  in  the  solid  and  gaseous  forms,  and  we  are 
chiefly  concerned  with  the  conditions  under  which  the  change  of  one 
form  into  the  other  takes  place,  the  phenomena  accompanying  these 
transformations,  and  the  conditions  under  which  two  or  more  forms 
of  a  substance  can  exist  together. 

The  Change  of  Liquid  to  Vapour.  Equilibrium— It  is  a 
well-known  fact  that  water  and  all  other  pure  liquids,  under  suitable 
conditions,  tend  to  give  off  vapour,  and  if  the  space  into  which  the 
vapour  escapes  is  sufficiently  large,  a  given  amount  of  a  liquid  may 
be  changed  completely  to  vapour.  The  process  of  transformation 
of  a  liquid  into  a  vapour  is  known  as  evaporation  or  vaporization. 

The  phenomenon  may  be  studied  most  advantageously  when 
vaporization  takes  place  into  a  confined  space.  Three  glass  tubes, 
A,  B,  and  C,  closed  at  one  end,  and  about  a  metre  long,  are  filled 
with  mercury  and  inverted  in  a  bath  of  the  same  metal  (Fig.  26). 
The  mercury  stands  at  the  same  level  in  each  tube.  So  much  water 
is  now  introduced,  by  means  of  a  bent  pipette,  into  the  upper  part  of 
the  tubes  B  and  C  that,  after  part  of  it  has  vaporized,  a  little  still 
remains  in  the  liquid  form.  The  vapour  exerts  a  pressure  which  is 
measured  by  the  difference  of  level  of  the  mercury  in  these  tubes  and 
in  the  comparison  tube  A.  If  one  of  the  tubes  is  lowered  a  little  in  the 
bath  so  as  to  diminish  the  space  above  the  mercury  some  vapour  at 
once  condenses ;  if  the  tube  is  raised  a  little  so  as  to  increase  the 
space  over  the  mercury  more  vapour  is  formed,  and  in  both  cases  the 
vapour  pressure  regains  the  original  value.  When  a  liquid  is  in 
contact  with  its  own  vapour  under  such  conditions,  the  space  is 
said  to  be  saturated  with  the  vapour.  It  follows  from  the  experi- 
ments just  described  that  at  constant  temperature  water  exerts  a 
definite  vapour  pressure,  which  can  be  reached  both  from  a  more 
saturated  and  from  a  less  saturated  state,  and  which  is  independent 
of  the  amount  of  liquid  present. 

The  above  results  are  best  expressed  by  the  statement  that  in  a 
confined  space  water  rapidly  attains  a  state  of  equilibrium  with  its 
vapour  ;  at  constant  temperature  the  space  above  the  liquid  contains 
a  definite  amount  of  vapour  per  unit  volume,  which  therefore  exerts  a 
definite  pressure  termed  the  vapour  pressure  of  the  liquid. 

The  vapour  pressure  of  all  liquids  increases  as  the  temperature 
rises.  If  the  tube  C  is  jacketed  and  heated  by  means  of  steam  from 
a  boiler,  it  will  be  noticed  that  the  mercury  is  more  and  more 


64    A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

depressed  as  the  temperature  rises,  and  finally,  when  the  whole  is 
heated  to  100°,  the  mercury  stands  at  the  same  level  outside  and 
inside  the  tube.  This  indicates  that  at  100°,  the  temperature  at 
which  water  boils,  its  vapour  pressure  is  equal  to  the  vapour  pressure 


FIG.  26. 

of  the  atmosphere.  The  same  is  true  for  all  other  liquids,  and  hence 
the  boiling-point  of  a  liquid  is  that  temperature  at  which  its  vapour 
pressure  is  equal  to  atmospheric  pressure.  In  the  following  table  the 
vapour  pressure  of  water  is  given  in  mms.  mercury  for  a  few  tempera- 
tures between  10°  and  1 50°  :— 


WATER— PHYSICAL   PROPERTIES    OF   LIQUIDS     65 


Temperature. 

Vapour  Pressure. 

Temperature. 

Vapour  Pressure. 

-10° 
+    0° 

+  10° 

+  20° 

+30° 

2.16  mm. 
4-58     „ 
9-!7     .. 
17.41     .. 
3i-56     •' 

+  40° 
+  50° 
+  70° 

+  100° 

+  150° 

54.97  mm. 
92.17    .1 
233'8     .. 
760.0    ,, 
358o.o     „ 

As  is  well  known,  the  boiling-points  of  liquids,  that  is,  the  tem- 
peratures at  which  they  exert  a  vapour  pressure  equal  to  that  of 
the  atmosphere,  are  very  different.  Thus,  of  the  pure  substances 
already  mentioned,  liquid  hydrogen  boils  at  —252.5°,  liquid  oxygen 
at  —  182.5°,  sulphur  dioxide  at  —  10°,  water  at  100°,  and  mercury 

at  356°. 

At  the  boiling-point  of  a  liquid  bubbles  of  vapour  form  and  rapidly 
escape  into  the  atmosphere,  which  gives  rise  to  the  characteristic 
bubbling  and  agitation  of  the  liquid  at  this  point.  It  is  evident  that 
a  liquid  will  boil  at  a  lower  temperature  when  the  pressure  above  it 
is  reduced — in  other  words,  when  the  pressure  of  the  air  is  reduced 
below  atmospheric  the  vapour  pressure  of  the  liquid  will  be  able  to 
overcome  it  at  a  lower  temperature.  This  may  readily  be  shown 
by  placing  a  flask  containing  water  at  40°  or  50°  in  connexion 
with  an  air-pump  and  rapidly  exhausting,  when  the  water  will  boil 
vigorously. 

The  equilibrium  between  liquid  and  vapour  may  also  be  considered 
from  the  standpoint  of  the  kinetic  theory.  Liquids,  like  gases,  may 
be  regarded  as  being  made  up  of  small  particles  in  more  or  less 
rapid  movement.  When  water  is  placed  on  the  top  of  mercury  in  a 
vacuum,  the  particles  in  most  rapid  motion  find  their  way  into  the 
space  above  the  liquid.  As  the  number  of  particles  in  the  vapour 
space  increases,  more  and  more  of  them  find  their  way  back  into 
the  liquid,  and  finally  a  condition  of  equilibrium  is  attained  in  which 
as  many  particles  enter  as  leave  the  liquid  in  a  given  time.  Accord- 
ing to  this  view,  the  equilibrium  between  a  liquid  and  its  vapour  is 
of  a  kinetic  and  not  of  a  static  character.  It  will  be  shown  later 
that  this  view  can  be  extended  to  chemical  as  well  as  to  physical 
equilibrium. 

The  amount  of  a  substance  in  unit  volume  is  conveniently  termed 

the  concentration  of  the  substance.    Thus  we  may  say  that  at  a  definite 

temperature  water  is  in  equilibrium  with  a  definite  concentration  of 

vapour,  and  the  higher  the  temperature  the  greater  is  the  concentra- 

5 


66     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

lion  of  vapour.  This  definition  of  concentration  should  be  carefully 
noted,  as  it  is  of  great  importance.  The  equilibrium  is  not  deter- 
mined by  the  absolute  amount  of  vapour,  but  by  the  amount  per  unit 
volume — the  concentration. 

In  the  foregoing,  we  have  considered  the  equilibrium  between 
liquid  and  vapour  in  the  absence  of  other  substances.  It  is  important 
to  remember,  however,  that  the  presence  of  another  vapour  or  gas 
does  not  affect  the  magnitude  of  the  vapour  pressure  of  a  liquid.  If, 
for  example,  one  of  the  tubes  represented  in  Fig.  26  contains  air,  the 
final  concentration  of  water  vapour  in  the  space  above  the  liquid  is 
the  same  as  when  water  vaporizes  into  a  vacuum.  The  only  effect 
of  the  presence  of  another  gas  or  vapour  is  that  the  equilibrium  is 
not  reached  so  rapidly. 

Heat  of  Vaporization— If  a  little  ether  is  poured  on  the  hand 
it  rapidly  evaporates,  and  a  powerful  cooling  effect  is  felt.  This  is  an 
illustration  of  the  fact  that  when  a  liquid  changes  to  a  vapour  heat 
is  absorbed.  The  heat  thus  taken  up  in  bringing  about  a  change  of 
state  is  termed  latent  heat ;  when  the  change  is  from  liquid  to 
vapour  it  is  called  latent  heat  of  vaporization.  It  has  already  been 
mentioned  that  the  latent  heat  of  vaporization  of  water  at  its  boiling- 
point  is  537  calories ;  in  other  words,  it  requires  the  expenditure  of 
537  calories  to  convert  I  gram  of  water  at  100°  to  water  vapour  at 
the  same  temperature.  The  same  amount  of  heat  is  of  course  given 
up  when  the  vapour  is  condensed.  We  can  now  understand  why  the 
temperature  of  boiling  water  remains  constant  although  heat  is  being 
continuously  supplied  ;  the  heat  is  used  up  in  bringing  about  a  change 
of  state. 

An  alternative  statement  of  the  facts  just  mentioned  is  that  steam 
has  much  more  energy  than  an  equal  weight  of  water.  According 
to  the  kinetic  theory,  the  heat  supplied  is  mainly  transformed  into 
kinetic  energy  ;  it  is  used  up  in  overcoming  the  attraction  between 
the  particles  (p.  49),  and  in  overcoming  the  pressure  of  the  atmosphere 
(p.  146). 

The  degree  of  cooling  which  may  be  attained  by  the  rapid  evapora- 
tion of  volatile  substances  is  very  considerable.  Thus  the  tem- 
perature of  liquid  ethylene,  which  boils  at  -  103°,  can  be  reduced  to 
-  120°,  and  in  the  same  way,  by  the  rapid  evaporation  of  liquid 
oxygen,  which  boils  at  -182.5°,  a  temperature  of  -210°  may  be 
reached. 

The  Change  of  Liquid  to  Solid— When  a  pure  liquid  is  pro- 
gressively cooled  it  finally  changes  to  the  solid  form,  and  conversely, 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     67 

when  the  solid  modification  is  heated  it  is  again  changed  to  the 
liquid  form  ;  in  other  words,  it  melts  or  fuses  at  a  definite  temperature. 
This  property  of  solidifying  or  melting  at  a  constant  temperature 
is  one  of  the  most  important  characteristics  of  pure  substances.  It 
may  be  mentioned  that  some  liquids  do  not  solidify  when  the  melting- 
point  is  reached,  but  may  require  to  be  supercooled  several  degrees 
before  the  solid  form  appears ;  when  solidification  once  begins,  how- 
ever, it  proceeds  at  constant  temperature,  which  is  the  same  as  that 
at  which  the  solid  fuses  (compare  next  section). 

In  the  case  of  water,  and  of  all  other  liquids  which  expand  on 
solidification,  the  temperature  at  which  solid  and  liquid  are  in  equi- 
librium is  lowered  by  increase  of  pressure,  but  the  effect  is  invariably 
small.  In  the  case  of  water  between  I  and  336  atmospheres  an 
increase  of  pressure  of  n  atmospheres  lowers  the  melting-point  of  ice 
by  0.0074  n  degrees,  so  that  ice  under  a  pressure  of  336  atmospheres 
melts  at  -2.5°.  At  higher  pressures  the  coefficient  is  somewhat 
greater.  Thus  under  a  pressure  of  1155  atmospheres  ice  melts  at 
-  10°,  and  under  a  pressure  of  2200  atmospheres  at  -  22°  (Tammann). 
The  melting-point  of  substances  which  contract  on  solidification  is 
raised  by  increase  of  pressure. 

Latent  Heat  of  Fusion— When  heat  is  supplied  to  a  solid  its 
temperature  rises  till  the  melting-point  is  reached,  and  then  remains 
constant  till  all  the  solid  is  melted.  The  phenomenon  is  exactly 
analogous  to  that  occurring  when  a  liquid  is  vaporized,  the  heat 
being  used  up  in  bringing  about  a  change  of  state.  Similarly,  when 
a  liquid  is  progressively  cooled,  the  temperature  falls  till  it  begins 
to  solidify,  and  then  remains  constant  till  solidification  is  complete, 
during  which  process  the  latent  heat  is  given  out. 

In  order  to  convert  I  gram  of  ice  at  o°  into  water  at  the  same 
temperature  80  calories  must  be  supplied  ;  in  other  words,  the  latent 
heat  of  fusion  of  ice  is  80  calories.  The  values  for  most  other  sub- 
stances are  lower  than  for  water.  Thus  the  latent  heat  of  fusion  of 
tin  is  14.25  calories  per  gram,  and  of  sulphur  9.37  calories  per  gram. 

Equilibrium  of  the  three  Modifications,  Ice,  Water,  and 
Vapour — The  proper  understanding  of  the  questions  discussed  in 
the  previous  paragraphs  is  greatly  facilitated  by  a  graphical  repre- 
sentation of  the  variation  of  physical  properties  with  temperature 
and  pressure.  The  accompanying  diagram  (Fig.  27)  affords  such  a 
representation  for  the  different  forms  of  water.  Two  lines  are  drawn 
at  right  angles,  the  so-called  rectangular  co-ordinates — and  tem- 
peratures are  measured  along  the  horizontal  axis,  pressures  along  the 


68     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Solid 


vertical  axis.  The  line  OA  represents  the  variation  of  the  vapour 
pressure  of  water  with  temperature  ;  it  is  obtained  by  drawing  a 
curve  through  the  points  at  which  lines  drawn  parallel  to  the  pressure 

axis  from  points  on  the  horizontal  axis 
representing  particular  temperatures 
meet  the  lines  drawn  parallel  to  the 
temperature  axis  from  points  on  the 
vertical  axis  representing  the  corre- 
sponding pressures.  The  point  O 
represents  the  temperature  at  which 
a  mixture  of  water  and  ice  is  in  equili- 
brium under  the  pressure  of  their  own 
vapour.  This  temperature  is  very 
near  to,  but  is  not  exactly  o°.  The 
latter  temperature  has  already  been 
defined  as  that  temperature  at  which 
ice  and  water  are  in  equilibrium  under 
atmospheric  pressure.  At  present, 


Vapour 


Temperature 

FIG.  27. 


however,  we  are  concerned  with  the  equilibrium  between  ice  and 
water  under  the  pressure  of  their  own  vapour.  As  this  amounts  at 
o°  to  about  4.6  mm.,  which  is  practically  an  atmosphere  below 
atmospheric  pressure,  and  as  diminution  of  pressure  raises  the 
melting-point  of  ice,  the  temperature  corresponding  with  the  point  O 
is  about  +  0.0074°. 

Like  water,  ice  exerts  a  definite  vapour  pressure,  as  is  evident  from 
the  fact  that  it  slowly  evaporates  at  temperatures  below  zero.  The 
line  or  curve  OC  represents  the  variation  of  the  vapour  pressure  of 
ice  with  temperature,  and  it  should  be  observed  that  it  is  not  a 
direct  continuation  of  the  curve  AO.  The  diagram  indicates  that  at 
o°  (strictly  speaking  at  +0.0074°)  water  and  ice  have  the  same  vapour 
pressure,  and  this  has  been  proved  both  experimentally  and  theoreti- 
cally. Suppose  that  water  and  ice  at  o°  are  contained  separately  in 
the  limbs  of  the  bent  tube  shown  in  Fig.  28,  and  let  us  assume  for 
a  moment  that  the  vapour  pressure  of  one  of  the  modifications,  say 
the  water,  is  greater  than  that  of  the  other.  Under  these  circum- 
stances the  water  would  continuously  pass  into  vapour,  which  would 
solidify  to  ice  in  the  other  limb,  until  all  the  water  has  disappeared. 
It  is  therefore  evident  that  the  two  modifications  can  only  be  in 
equilibrium  when  they  exert  the  same  vapour  pressure. l 

As  water  can  be  supercooled  considerably,  its  vapour  pressure  can 

1  The  change  of  a  solid  into  vapour  and  recondensation  of  the  latter  to  solid  is 
termed  sublimation. 


WATER— PHYSICAL   PROPERTIES    OF    LIQUIDS     69 

be  measured  for  some  degrees  below  zero.  The  vapour  pressure  of 
supercooled  water  is  represented  by  the  line  OA'  on  the  diagram, 
which  lies  above  OC  ;  it  follows,  therefore,  that  the  vapour  pres- 
sure of  supercooled  water  is  greater  than  that  of  ice  at* the  sam& 
temperature.  This  at  once  explains  why  supercooled  water  cannot 
exist  in  contact  with  ice ;  if  they  are  contained  in  the  bent  tube 
(Fig.  28)  at  a  temperature  below  o°,  the  water  will  completely 
evaporate  into  the  other  limb  in  virtue  of  its  higher  vapour  pressure 
and  solidify. 

Supercooled  water  is  sometimes  said  to  be  unstable,  as  it  solidifies 
at  once  in  contact  with  ice.  It  is  preferable  to  use  the  term  meta- 
stable  in  this  connexion  in  order  to  indicate  that  supercooled  water 
has  little  or  no  tendency  to  crystallize  in  the  entire  absence  of  the 
solid  form.  These  statements  apply  to  other  substances  as  well ; 
the  metastable  form  of  a  substance  has  invariably  a  higher  vapour 
pressure  than  the  stable  form.  It  will,  of  course,  be  understood  that 
relative  stability  or  instability  is  entirely  a  question  of  conditions. 
Water  is  stable  above  o°,  metastable  below  o°,  ice  stable  below  o°,  and 
would  be  metastable  above  o°  if  it  could  be  superheated  (cf.  p.  85). 

The  equilibrium  diagram  for  water  may  be  completed  for  our 
present  purpose  by  drawing  the  line  OB,  which  represents  the  effect 
of  pressure  on  the  melting-point  of  ice.  As  increased  pressure|lowers 
the  melting-point,  the  line  is  slightly  inclined  towards  the  pressure 
axis,  as  shown. 

It  is  evident  that  at  points  on  the  curves  only  two  forms  are  in 
equilibrium,  solid  and  liquid  along  OB,  liquid  and  vapour  along  OA, 
and  ice  and  vapour  along  OC.  At  only  one  point,  the  point  O,  are 
three  modifications  in  equilibrium,  and  the  point  O  is  therefore  termed 
a  triple  point.  The  temperature  (0.0074°)  and  pressure  (4.6  mm.)  at 
the  triple  point  are  the  only  values  at  which  the  three  forms  can. 
exist  together.  If  the  pressure  is  increased,  vapour  disappears ;  if 
it  is  diminished,  water  disappears.  If  the  temperature  is  increased, 
ice  disappears  ;  if  it  is  lowered,  water  disappears.  As  the  diagram 
shows,  at  the  values  of  temperature  and  pressure  represented  by  the 
regions  between  the  curves  only  one  form  is  capable  of  existence. 

Delay  in  Appearance  of  Itfew  Forms.  Phases— Water, 
ice,  and  water  vapour  are  often  termed  different  phases  of  the  same 
substance.  Each  phase  is  homogeneous  throughout  and  is  separated 
by  a  definite  surface  from  other  phases.  The  term  is  chiefly  used  when 
one  is  dealing  with  so-called  heterogeneous  systems,  which  are  made 
up  of  two  or  more  phases.  Water  in  contact  with  water  vapour  is 
such  a  heterogeneous  system,  made  up  of  two  phases,  a  liquid  and 


70     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

a  vapour  phase.  This  term  will  be  found  very  useful  in  our  latel 
work. 

We  have  already  learnt  that  ice  does  not  necessarily  appear  when 
water  is  'carefully  cooled  below  o°.  This  and  similar  facts  may  be 
generalized  in  the  statement  that  when  the  conditions  are  such  that 
a  new  phase  may  appear,  it  does  not  necessarily  form  unless  a  small 
amount  of  it  is  already  present  in  the  system.  We  shall  see  later 
that  the  amount  of  a  new  phase  necessary  to  ensure  the  appearance 
of  the  latter  in  quantity  under  favourable  conditions  is  exceedingly 
small  (cf.  p.  85). 

A  further  illustration  of  retardation  in  the  appearance  of  a  new 
phase  is  that  water  may  be  heated  in  a  clean  vessel  several  degrees 
above  its  boiling-point  without  the  formation  of  vapour.  At  a  certain 
point,  however,  a  considerable  quantity  of  vapour  suddenly  escapes, 
thus  giving  rise  to  that  irregular  form  of  boiling  known  as  "bumping." 

In  the  absence  of  liquid  water  (and  of  dust),  water  vapour  can  be 
obtained  under  pressures  greater  than  its  vapour  pressure  under  the 
experimental  conditions  without  the  appearance  of  water.  On  the 
other  hand,  as  already  mentioned,  ice  cannot  be  superheated  ;  the 
new  phase  appears  as  soon  as  the  temperature  exceeds  o°. 

Liquefaction  of  Gases— We  have  seen  in  Chapter  V.  that  a 
liquid  is  in  equilibrium  with  its  own  vapour  at  a  definite  temperature 
and  pressure.  If  at  constant  pressure  the  temperature  is  raised  and 
kept  at  the  new  value  the  whole  of  the  liquid  will  vaporize  ;  if  under 
the  same  circumstances  the  temperature  is  kept  below  the  equilibrium 
value,  the  whole  of  the  vapour  will  liquefy.  Similarly,  at  constant 
temperature  the  formation  of  liquid  is  favoured  by  increasing,  the 
formation  of  vapour  by  lowering,  the  pressure. 

From  this  it  would  appear  to  follow  that  any  gas  or  vapour  can  be 
liquefied  by  sufficiently  lowering  the  temperature  and  increasing  the 
pressure,  and,  as  a  matter  of  fact,  every  known  gas  has  now  been 
liquefied  by  applying  this  principle.  One  important  point  must, 
however,  be  noted  in  this  connexion.  If  carbon  dioxide  is  confined 
in  a  tube  at  room  temperature  and  the  pressure  gradually  increased, 
at  a  certain  point  liquid  will  appear  in  the  tube,  and  by  further 
increasing  the  pressure  the  whole  of  the  gas  liquefies.  Under  the 
same  circumstances,  however,  no  pressure,  however  great,  brings 
about  the  liquefaction  of  oxygen.  These  facts  puzzled  the  older 
chemists,  but  the  problem  was  finally  solved  by  Andrews,  who 
showed  that  for  every  gas  or  vapour  there  is  a  temperature  (which 
differs  for  each  gas)  below  which  it  can  be  liquefied  by  pressure,  but 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     71 


above  which  no  amount  of  pressure  can  bring  about  transformation 
to  the  liquid  state.  This  temperature  is  called  the  critical  tempera- 
ture. It  is  now  easy  to  understand  the  different  behaviour  of  carbon 
dioxide  and  of  oxygen  under  pressure.  The  critical  temperature  of 
carbon  dioxide  is  31°,  so  that  at  room  temperature  it  is  below  its 
critical  temperature  ;  on  the  other  hand,  the  corresponding  value  for 
oxygen  is  —  119°,  so  that  at  room  temperature  it  is  far  above  its  critical 
temperature. 

The  pressure  just  sufficient  to  liquefy  a  gas  at  the  critical  tempera- 
ture is  termed  the  critical  pressure.  For  carbon  dioxide  at  31°  the 
critical  pressure  is  72  atmospheres.  For  oxygen  at -119°  it  is 
58  atmospheres.  It  will,  of  course,  be  readily  understood  that  the 
further  a  gas  or  vapour  is  cooled  below  the  critical  temperature  the 
smaller  is  the  pressure  required  to  liquefy  it. 

The  physical  constants,  including  the  critical  temperatures  and 
pressures,  of  some  of  the  commoner  gases  is  given  in  the  accom- 
panying table. 


Boiling- 
point. 

Melting- 
point. 

Critical 
Tempera- 
ture. 

Critical 
Pressure. 

Density  at 
Boiling- 
point. 

0 

0 

0 

Atmos. 

Helium     . 

-268.5 

-268? 

2-75 

0.15 

Hydrogen 
Nitrogen  . 

-252.5 
-195.6 

-259 
-213 

-241 

-149 

14 
27-5 

0.07 
0.791 

Carbon  monoxide 

-190 

-207 

-136 

33-4 

Oxygen     . 

-182.5 

-223 

-119 

58.0 

1.131 

Ethylene  . 

-i°3-5 

-169 

+     9 

58.0 

0-571 

Nitrous  oxide   . 

-  89.8 

-102.7 

+  37 

73-o 

1.226 

Carbon  dioxide 

-  80 

... 

+  31-35 

72.3 

... 

Ammonia  . 

-   38.5 

-  75-5 

+  131 

"3 

... 

Chlorine    . 

-   33-4 

-102 

+  141 

84 

I-507 

Sulphur  dioxide 

-     10 

-    76 

+  I5S-4 

79 

4 

Methods  used  in  Liquefying  Gases— The  first  example  of 
the  conversion  of  a  substance  which 
is  a  gas  at:  the  ordinary  temperature 
into  a  liquid  by  pressure  was  chlorine 
(Northmore,  1806).  Later,  Faraday 
succeeded  in  liquefying  a  number  of 
the  commoner  gases,  such  as  sulphur 
dioxide,  nitrous  oxide,  and  ammonia.  For  this  purpose  he  used  bent 
tubes  of  strong  glass  (Fig.  28) ;  substances  for  generating  the  gas  were 


FIG.  28. 


72     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

placed  in  one  limb  <z,  the  end  b  of  the  tube  was  then  sealed  up,  and 
on  heating  the  other  end  the  gas  was  given  off,  and  under  its  own 
pressure  part  of  it  condensed  in  the  limb  b.  If  necessary,  this  end 
could  be  placed  in  a  freezing  mixture. 

Faraday  did  not  succeed  in  liquefying  oxygen  or  air  ;  the  credit  of 
having  first  obtained  oxygen  as  a  coherent  liquid  is  due  to  Pictet 
(1877).  The  gas  was  obtained  under  very  great  pressure  by  heating 
potassium  chlorate  in  a  confined  space,  while  the  copper"  tube  con- 
taining the  gas  was  surrounded  by  liquid  carbon  dioxide  kept  at  -  120° 
10—140°  by  rapid  evaporation  under  reduced  pressure  (p.  66).  The 
liquefied  carbon  dioxide  used  in  this  experiment  was  obtained  by 
compressing  the  gas  in  a  tube  surrounded  by  liquefied  sulphur  dioxide 
evaporating  under  reduced  pressure  (temperature  — 60°) ;  the  liquid 
sulphur  dioxide  in  its  turn  being  obtained  by  compressing  the  gas  at 
room  temperature.  Prior  to  the  working  out  by  Linde  and  Hampson 
(independently)  of  the  modern  method  of  liquefying  air  and  other  so- 
called  "permanent"  gases,  Pictet's  method  was  exclusively  used  for 
this  purpose.  Simultaneously  with  Pictet,  Cailletet  obtained  liquid 
oxygen,  though  only  in  the  form  of  a  mist  and  in  small  drops,  by 
subjecting  the  gas  to  great  pressure,  which  was  then  suddenly 
released.  The  intense  cooling  thus  obtained  caused  the  momentary 
liquefaction  of  part  of  the  compressed  gas. 

The  older  methods  have  now  been  completely  displaced,  as  regards 
the  less  condensible  gases,  by  the  Linde-Hampson  method,  which 
has  led  to  the  liquefaction  of  hydrogen  and  helium.  The  principle  of 
the  method  is  that  when  a  gas  is  allowed  to  pass  from  a  high  to  a 
low  pressure  through  a  small  opening  (jet)  or  series  of  small  openings 
(porous  plug)  it  becomes  cooled  (Joule-Thomson  effect).  The  cooling 
effect  is  due  to  the  performance  of  "  internal "  work  in  overcoming 
the  mutual  attraction  of  the  particles,  and  is  therefore  observed  only 
for  "imperfect"  gases  (p.  50).  The  effect  is  the  greater  the  lower 
the  temperature  at  which  the  expansion  takes  place,  and  the  greater 
the  difference  of  pressure  on  the  two  sides  of  the  valve.  The  cooling 
effects  thus  obtained  are  summed  up  in  a  very  ingenious  way  by  the 
principle  of  "contrary  currents,"  the  same  quantity  of  gas  being 
caused  to  circulate  through  the  apparatus  several  times  ;  after  expand- 
ing through  the  jet  it  is  caused  to  flow  over  and  cool  the  tube  through 
which  a  further  quantity  of  gas  is  passing  on  its  way  to  the  jet. 

The  apparatus  employed  is  represented  diagramatically  in  Fig.  29. 
By  means  of  the  pump  A  the  gas  is  compressed  in  B  to  (say)  100 
atmospheres,  the  heat  given  out  in  the  process  being  absorbed  by 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     73 


surrounding  B  with  a  vessel  through  which  a  current  of  cold  water  is 
passed.  The  cooled,  compressed  gas  then  passes  down  the  central 
tube  G  towards  the  jet  E,  being  further  cooled  on  the  way  by  the  gas 
passing  up  the  wide  tube  D,  which  has 
just  expanded  through  the  jet.  After 
passing  through  E,  and  thus  falling  to 
its  original  pressure,  the  gas  passes  up- 
wards over  the  central  tube  G  and  again 
reaches  A  by  the  tube  C  and  the  left- 
hand  valve  at  the  bottom  of  A.  The 
direction  of  the  circulating  stream  of 
gas  is  indicated  by  the  arrows.  In 
course  of  time  the  temperature  becomes 
so  low  that  part  of  the  gas  is  liquefied,  c 
and  collects  in  the  vessel  F.  More  air 
is  drawn  into  the  apparatus  as  required, 
and  the  process  is  continuous. 

By  means  of  an  apparatus  constructed 
on  this  principle,  Dewar,  and,  some- 
what later,  Travers,  succeeded  in  ob- 
taining liquid  hydrogen  in  quantity. 
The  hydrogen  before  expanding 
through  the  jet  was  cooled  to  about 
—  200°  by  means  of  liquid  air  boiling 
under  reduced  pressure.  In  the  same 
way  Kammerlingh  Onnes  of  Leyden 
has  just  succeeded  in  liquefying  helium, 
the  preliminary  cooling  in  this  case 
being  effected  by  means  of  liquid 
hydrogen.  All  known  gases  have  now 
been  liquefied. 

The  Linde  -  Hampson  machine  is 
chiefly  used  for  making  liquid  air, 
which  is  now  a  relatively  cheap  com- 
mercial article.  As  will  be  shown  in 
detail  later,  air  is  essentially  a  mixture 


FIG.  29. 


of  the  two  "permanent"  gases  oxygen  and  nitrogen.  The  boiling- 
point  of  liquid  air  lies  between  those  of  the  two  components,  oxygen 
- 182.5°  and  nitrogen  195.6°,  and  when  the  liquid  is  fresh  is  about 
-  190°.  It  contains  about  50  per  cent,  of  oxygen.  As  nitrogen  boils 
at  a  lower  temperature  than  oxygen,  the  former  passes  off  in  much 


74    A  TEXT-BOOK  OF   INORGANIC  CHEMISTRY 

greater  proportion  as  the  liquid  boils,  and  finally  a  mixture  very  rich 
in  oxygen  is  obtained.  As  already  mentioned,  this  method  is  used 
for  the  commercial  preparation  of  oxygen  (p.  24).  The  first  fractions, 
which  consist  of  almost  pure  nitrogen,  are  also  of  commercial  value 
(p.  236).  This  process  of  separating  a  liquid  into  its  components  by 
taking  advantage  of  a  difference  in  their  boiling-points,  is  known  as 
fractional  distillation. 

A  much  more  efficient  separation  of  the  constituents  of  air  is 
secured  by  methods  devised  by  Claude  and  by  Linde,  which  are 
based  on  a  process  which  has  long  been  in  use  for  extracting  spirit 
from  a  weak  solution  of  alcohol  and  water.  The  weak  alcohol  is 
made  to  trickle  down  a  tower  containing  zig-zag  shelves  or  baffle 
plates,  and  a  current  of  steam  is  admitted  at  the  bottom  and  passed 
up  through  the  shelves.  At  each  stage  some  of  the  alcohol  (which 
boils  at  a  lower  temperature  than  water)  is  vaporized,  and  some 
of  the  steam  condensed,  and  finally  a  vapour  very  rich  in  alcohol 
escapes  at  the  top  and  almost  pure  water  trickles  out  below.  In  the 
application  of  this  method  to  liquid  air,  the  nitrogen,  as  the  more 
volatile  substance,  corresponds  with  the  alcohol  and  the  oxygen 
with  the  water.  Liquid  air  trickles  down  through  a  rectifying  column 
up  which  nearly  pure  oxygen  (obtained  by  evaporation  of  liquid  air) 
is  passed.  As  the  gas  passes  up  the  column  there  is  a  continuous 
exchange  of  substance  ;  at  each  stage  some  of  the  rising  oxygen  is 
condensed,  and  some  of  the  nitrogen  in  the  downcoming  liquid 
is  evaporated  ;  finally  almost  pure  oxygen  is  obtained  at  the  bottom, 
and  nitrogen  containing  a  comparatively  small  proportion  of  oxygen 
passes  off  at  the  top.  Quite  recently,  Claude  has  still  further  im- 
proved the  process  and  obtained  a  complete  separation  of  the  oxygen 
and  nitrogen. 

Claude's  Method  of  Preparing  Liquid  Air1— It  has  been  pointed  out  that 
in  the  Linde  method  of  preparing  liquefied  air  no  external  work  is  done ;  the 
cooling  is  done  by  internal  work.  Many  attempts  have  been  made  to  liquefy 
air  by  allowing  it  to  expand  in  a  working  cylinder  with  performance  of  external 
work  (p.  146),  but  at  first  the  question  of  lubrication  at  such  low  temperatures 
presented  a  serious  difficulty.  Suitable  lubricants  have  now  been  found  by 
Claude  in  petroleum  ether,  which  becomes  viscous,  but  does  not  solidify  even  at 
- 160° ;  at  still  lower  temperatures  liquid  air  itself  is  used.  In  the  first  stage 


1  An  excellent  account  of  modern  methods  of  liquefying  air  and  other  gases  is 
given  by  Ewing — The  Mechanical  Production  of  Cold  (Cambridge  University 
Press). 


WATER— PHYSICAL   PROPERTIES   OF   LIQUIDS     75 

of  Claude's  process,  air  is  cooled  by  expansion  in  a  working  cylinder  to  a 
temperature  rather  below  - 140°  (the  critical  tem- 
perature of  air) ;  this  cooled  air  is  then  used  to  cool 
the  remainder  of  the  air,  with  the  result  that  the 
latter  is  liquefied  at  the  pressure  (50  atmospheres)  at 
which  it  is  supplied  to  the  apparatus.  It  is  doubtful 
whether  the  preparation  of  liquid  air  by  external 
expansion  alone  is  more  advantageous  than  the 
Linde  method,  but  the  use  of  expansion  with  external 
work  for  the  first  stage  of  the  cooling,  the  cooled  air 
being  then  liquefied  by  expansion  through  a  throttle 
valve  (without  external  work)  appears  to  present  un- 
doubted advantages. 

Investigations  at  low  temperatures  have  been 
greatly  facilitated  by  the  use  of  so-called  Dewar 
flasks  —  glass  vessels  with  double  walls,  the  space 
between  the  walls  being  completely  exhausted  of 
air  (Fig.  30).  Since  heat  is  conveyed  much  more 
slowly  through  a  vacuum  than  through  a  space  con- 
taining a  gas,  the  conduction  of  heat  from  the  outside 
is  reduced  to  such  an  extent  that  liquid  air  can  be 


FIG.  30. 


preserved  for  hours  with  very  little  loss.  A  still  better  result  is  obtained  by 
coating  the  walls  with  silver;  by  this  means  radiant  heat  is  reflected.  If  the 
vessel  is  plugged  with  glass  wool,  the  air  confined  in  the  pores  acts  as  a  very 
efficient  non-conductor  and  the  evaporation  is  still  further  diminished. 

In  recent  years,  owing  to  their  slight  conducting  power  for  heat,  Dewar  vessels 
have  come  largely  into  use  as  calorimeter  vessels  in  accurate  work. 


CHAPTER  VII 
SOLUTION 

IN  the  preceding  chapters,  frequent  reference  has  been  made  to 
solutions,  more  particularly  those  of  gases  and  of  solids  in  liquids, 
and  it  will  be  convenient  to  summarize  at  this  stage  some  of  the  more 
important  phenomena  of  solution,  illustrated  mainly  by  the  examples 
already  quoted. 

When  a  little  sugar  is  shaken  up  with  water,  the  former  quickly 
disappears  and  a  perfectly  homogeneous  mixture  is  obtained,  which 
has  a  sweet  taste,  but  in  which  no  particles  of  sugar  can  be  detected, 
even  under  the  highest  power  of  the  microscope.  The  sugar  does 
not  settle  out  of  the  water,  no  matter  how  long  the  mixture  may  be 
kept,  but  can  be  obtained  in  its  original  form  by  evaporating  off  the 
water.  If,  on  the  other  hand,  a  little  chalk  is  shaken  up  with  water, 
a  milky  mixture  is  obtained  in  which'  the  particles  of  chalk  can 
readily  be  detected  under  the  microscope,  and  on  standing  the  chalk 
sinks  to  the  bottom,  leaving  the  water  clear.  In  the  former  case  the 
sugar  is  taken  up  or  dissolved  by  the  water,  forming  a  true  solution, 
in  the  latter  case  the  chalk  is  merely  suspended  in  the  water.  In 
these  extreme  cases  there  is  no  difficulty  in  distinguishing  between 
a  suspension  and  a  true  solution,  but  we  shall  see  later  that  some- 
times the  distinction  is  by  no  means  easy  to  draw. 

A  solution  may  be  defined  as  a  homogeneous  mixture  of  two  or 
more  substances,  and  to  this  we  may  add  that  the  composition  of  a 
solution  can  be  'varied  continuously  within  certain  limits.  The  latter 
part  of  this  definition  serves  to  distinguish  between  a  solution  and  a 
chemical  compound.  The  latter,  as  we  have  already  learnt,  are  of 
definite  and  invariable  composition,  whereas  the  composition  of  a 
solution  of  sugar  in  water  may  vary  within  wide  limits. 

The  component  of  a  solution  which  is  present  in  largest  propor- 
tion is  often  called  the  solvent,  and  the  substance  taken  up  by  the 
solvent  is  termed  the  dissolved  substance  or  solute.  In  some  cases, 
however,  as  in  a  mixture  of  alcohol  and  water  in  equal  volumes,  it 

cannot  be  said  that  one  of  the  compounds  has  more  right  than  the 

76 


SOLUTION  77 

other  to  be  regarded  as  the  solvent.  Strictly  speaking,  therefore,  no 
sharp  distinction  can  be  drawn  between  the  terms  solvent  and  solute, 
but  as  a  matter  of  convenience  they  are  often  employed  in  the  sense 
already  indicated. 

The  extent  to  which  one  substance  can  take  up  another  to  form  a 
solution  depends  largely  on  the  nature  of  the  substances.  Thus 
water  and  alcohol  are  miscible  in  all  proportions,  whilst  chalk  is 
scarcely  taken  up  at  all  by  water.  When  the  miscibility  is  limited, 
and  a  solvent  has  taken  up  as  much  of  a  solute  as  it  can  retain  in 
contact  with  the  undissolved  solute  (compare  p.  84),  the  solution 
is  said  to  be  saturated.  A  saturated  solution  represents  a  state  of 
equilibrium  between  two  phases,  ijie  solution  and  the  undissolved 
substance. 

Solutions  are  usually  classified  according  to  the  physical  state  of 
the  components.  The  more  important  classes  are  as  follows  : — 

(1)  Solutions  of  gases  in  gases, 

(2)  Solutions  of  gases  in  liquids, 

(3)  Solutions  of  liquids  in  liquids, 

(4)  Solutions  of  solids  in  liquids, 

and  each  of  these  will  be  briefly  considered. 

Solutions  of  Gases  in  Gases  — Gases  are  miscible  in  all 
proportions,  in  other  words,  the  mutual  solubility  of  gases  is  unlimited. 
An  important  law  which  has  been  established  as  the  result  of  the 
investigation  of  mixtures  of  gases  is  that  when  two  gases,  which  do 
not  act  chemically  on  each  other,  are  mixed  the  total  pressure  is  the 
sum  of  the  pressures  of  the  two  gases  taken  separately.  Otherwise 
expressed,  the  pressure  exerted  by  each  of  the  gases  is  the  same  as 
if  the  other  gas  were  not  present,  and  the  same  mass  of  any  one  of 
the  gases  occupied  the  total  space.  The  pressure  exerted  by  any 
one  gas  in  a  mixture  of  gases  is  termed  its  partial  pressure.  The 
fact  that  the  vapour  pressure  of  water  is  the  same  in  air  as  in  a 
vacuum  under  equivalent  conditions  (p.  66)  is  an  excellent  illustra- 
tion of  the  law.  The  above  law,  which  was  discovered  by  Dalton, 
is  a  particular  case  of  a  more  general  law,  which  states  that  when 
the  components  of  a  gaseous  mixture  exert  no  mutual  influence,  the 
properties  of  the  mixture  are  the  sum  of  the  properties  of  the  con- 
stituents. This  law  is  of  the  same  order  of  validity  as  the  gas  laws 
already  enumerated  (p.  40),  and  is  the  more  nearly  true  the  smaller 
the  concentrations  of  the  gaseous  components. 

Solutions  of  Gases  in  Liquids  —  Unlike  the  solubility  of 
gases  in  gases,  the  solubility  of  gases  in  liquids  is  limited.  The 


78     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


amount  of  a  gas  which  a  liquid  can  take  up  depends  upon  four 
factors  :  (a)  the  nature  of  the  gas,  (b)  the  nature  of  the  liquid,  (c)  the 
temperature,  (d)  the  pressure. 

(a)  The  amounts  of  different  gases  which  the  same  solvent  can 
take  up  vary  enormously  with  the  nature  of  the  gas.  Thus  i  c.c.  of 
water  at  o°  absorbs  0.0203  c.c.  of  hydrogen,  1.713  c.c.  of  carbon 
dioxide,  and  1148  c.c.  of  ammonia  gas  under  76  cm.  pressure. 

(ft)  The  amounts  of  the  same  gas  taken  up  by  different  solvents 
under  comparable  conditions  vary  greatly  with  the  nature  of  the 
solvent.  The  volumes  of  hydrogen  taken  up  by  I  c.c.  of  each  of 
the  following  solvents  at  o°  under  a  gas  pressure  of  76  cm.  is  as 
follows  : — 


Solvent 
Solubility"  . 


Water.      Aniline. 


Acetic 
Acid. 


0.0203       0.0303       0.06l7        0.0764       0.08SI       0.0862 


(c)  The  solubility  of  gases  diminishes  regularly  as  the  temperature 
rises.  This  rule  is  illustrated  by  the  following  table,  which  gives  the 
volumes  of  hydrogen,  nitrogen,  and  carbon  dioxide  taken  up  by  I  c.c. 
of  water  at  temperatures  between  o°  and  50°  under  a  gas  pressure  of 
76  cms. : — 


Solute. 

"  Absorption  Coefficient  "  in  Water  at 

0° 

10° 

20° 

30° 

50° 

Hydrogen   .        . 
Nitrogen     . 

0.0203 
0.0239 

0.0190 
0.0196 

0.0177 
0.0164 

0.0163 
0.0138 

0.0146 

0.0106 

Carbon  dioxide  . 

I-7I34 

1.194 

0.878 

0.665 

0.436 

The  majority  of  gases  can  be  completely  removed  from  solution 
by  raising  the  temperature.  This,  however,  is  not  invariably  the 
case  ;  hydrogen  chloride,  for  instance,  cannot  be  completely  removed 
from  aqueous  solution  by  boiling  (cf.  p.  95). 

(d)  The  amount  of  a  gas  absorbed  by  a  given  volume  of  liquid 
increases  with  increased  pressure.  The  law  connecting  the  amount 
of  absorption  and  pressure,  usually  known  as  Henry's  law,  is  as 
follows  :  At  constant  temperature,  the  amount  of  gas  absorbed  by  a 
given  volume  of  liquid  is  proportional  to  the  pressure  of  the  gas. 
Another  way  of  stating  Henry's  law  is  that  the  volume  of  gas  taken 
up  by  a  given  volume  of  liquid  is  independent  of  the  pressure.  This 
is  clearly  equivalent  to  the  first  statement,  because  when  the  pressure 


SOLUTION  79 

is  doubled  the  quantity  of  gas  absorbed  is  doubled,  but  since  its 
volume,  according  to  Boyle's  law,  is  halved,  the  original  and  final 
volumes  absorbed  are  the  same.  A  third  instructive  method  of  stating 
Henry's  law  is  that  the  ratio  of  the  concentration  of  the  dissolved  gas 
to  that  in  the  free  space  above  the  liquid  is  independent  of  the  pressure. 
In  the  case  of  hydrogen  at  o°,  for  example,  the  ratio  of  the  concen- 
tration in  water  to  that  in  the  gas  space  is  0.0203  : 1,  or  approximately 
i  :  50.  When  the  pressure  is  doubled,  the  concentration  both  in  the 
gas  space  and  in  water  is  doubled,  but  the  ratio  remains  unaltered. 

Henry's  law  is  approximately  valid  for  such  gases  as  hydrogen  and 
oxygen,  but  does  not  hold  for  very  soluble  gases  nor  for  those  gases 
which  enter  into  chemical  combination  with  the  solvent  (p.  95). 

For  comparative  purposes,  the  solvent  power  of  a  liquid  for  a  gas 
is  best  expressed  in  terms  of  the  "solubility"  or  "coefficient  of 
solubility,"  which  is  the  volume  of  the  gas  taken  up  by  unit  volume 
of  the  liquid  at  a  definite  temperature  (Ostwald).  As  is  clear  from 
the  foregoing,  the  "solubility"  is  simply  the  ratio  in  which  the  gas 
distributes  itself  in  the  liquid  and  in  the  gas  space,  and  has  the  great 
advantage  of  being  independent  of  the  pressure  in  so  far  as  the  gas 
follows  Henry's  law. 

Solubility  measurements  are  still  sometimes  expressed  in  terms  of 
the  so-called  "  absorption  coefficient "  of  Bunsen,  which  is  the 
volume  of  gas,  reduced  to  o°  and  76  cm.  pressure,  absorbed  by  i  c.c. 
of  a  liquid  at  a  definite  temperature  under  a  gas  pressure  equal  to  76 
cm.  of  mercury.  A  little  consideration  shows  that  the  only  difference 
between  the  "solubility"  s  and  the  "absorption  coefficient"  a,  is 
that  in  the  former  case  the  volume  of  the  gas  is  taken  at  the  tem- 
perature of  the  experiment  and  in  the  latter  case  is  reduced  to  o°. 
For  measurements  at  o°,  therefore,  the  two  factors  coincide,  and  at 
other  temperatures  the  relationship  is  expressed  by  the  equation 
s/a  =  (273  +  ^/273,  where  t  is  the  temperature  of  observation  on  the 
Centigrade  scale. 

A  convenient  form  of  apparatus  for  determining  the  solubility  of 
gases  in  liquids,  used  in  Ostwald's  laboratory,  is  represented  in 
Fig.  31.  The  graduated  measuring  tube  A  is  connected  by  a  rubber 
tube  with  the  tube  B,  and  by  altering  the  height  of  the  latter  till 
the  liquid  (mercury  or  water)  stands  at  the  same  height  in  the  tubes, 
a  quantity  of  gas  confined  in  the  upper  portion  of  A  can  readily  be 
measured  at  atmospheric  pressure.  The  tube  A  is  connected  with 
the  absorption  vessel  C  by  a  flexible  metal  tube;  a  and  b  are  three- 
way  stopcocks  and  c  is  a  simple  stopcock.  In  making  an  experiment 


8o     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


the  vessel  C,  the  volume  of  which  must  be  known,  is  filled  with  the 
air-free  solvent  and  a  quantity  of  the  gas  is  brought  into  A.  The 

volume  of  the  gas  at  atmospheric 
pressure  is  then  noted,  A  and  C  are 
put  in  connexion  by  means  of  the 
stopcocks  a  and  b,  and  a  measured 
volume  of  liquid  run  out  of  C,  its 
place  being  taken  by  the  gas.  The 
tap  b  is  then  closed,  the  gas  and 
liquid  in  C  are  vigorously  shaken  in 
order  to  secure  equilibrium,  the  tap 
b  again  opened  and  the  diminution 
in  the  volume  of  the  gas  read  off  in 
A  after  adjustment  to  atmospheric 
pressure.  As  the  volume  of  liquid 
in  C  'and  the  volume  of  gas  absorbed 
are  known,  the  solubility  of  the  gas 
at  the  temperature  of  the  experiment 
can  readily  be  calculated.  In  order 
to  secure  constancy  of  temperature, 
A  and  C  are  immersed  in  a  bath  of 
water  or  are  surrounded  by  vapours 
of  known  temperature. 

Solubility  of  Mixed  Gases- 
It  has  already  been  pointed  out  that 
each  component  of  a  gaseous  mixture 

exerts  its  effect  quite  independently  of  the  other  components.  The 
solubility  of  mixed  gases  is  a  particular  case  of  this  law ;  the  amount 
of  any  one  gas  taken  up  by  a  given  volume  of  liquid  is  proportional 
to  its  partial  pressure  in  the  mixture.  This  law  was  established  by 
Dalton  (1807),  and  is  known  as  Dalton's  law  of  partial  pressures. 

In  calculating  the  amounts  of  each  gas  taken  up  from  a  mixture, 
the  solubility  of  the  gas,  as  well  as  its  partial  pressure,  must  of  course 
be  taken  into  account.  Thus  if  a  mixture  of  two  volumes  of  hydrogen 
and  one  volume  of  nitrogen,  which  have  about  the  same  solubility  in 
water  at  o°,  is  shaken  up  with  water,  the  volume  of  hydrogen  dissolved 
will  be  about  double  that  of  the  nitrogen,  but  if  a  mixture  of  two 
volumes  of  hydrogen  and  one  volume  of  oxygen  are  used,  approxi- 
mately equal  volumes  of  the  gases  will  be  dissolved,  as  the  coefficient 
of  solubility  of  oxygen  is  about  double  that  of  hydrogen. 
Solutions  of  Liquids  in  Liquids  — As  regards  the  mutual 


FIG.  31. 


SOLUTION  81 

solubility  of  liquids  three  cases  may  be  distinguished:  (i)  The 
liquids  mix  in  all  proportions,  e.g.  alcohol  and  water;  (2)  the  liquids 
are  practically  immiscible,  e.g.  benzene  and  water  ;  (3)  the  liquids  are 
partially  miscible,  e.g.  ether  and  water. 

In  some  respects  the  partially  miscible  liquids  are  the  most  in- 
teresting. The  mutual  solubility  of  ether  and  water  may  be  shown 
by  shaking  approximately  equal  volumes  of  the  two  liquids  in  a 
separating  funnel.  On  allowing  to  stand,  a  separation  into  two  layers 
takes  place,  the  upper  layer  consisting  of  a  solution  of  water  in  ether, 
and  the  lower  of  a  solution  of  ether  in  water.  The  lower  layer  may 
be  separated  from  the  upper  by  carefully  opening  the  tap,  and  the 
presence  of  ether  may  be  shown  by  warming  the  solution  in  a  flask 
provided  with  a  cork  and  glass  tube  ;  the  escaping  ether  can  be 
•  ignited  at  the  end  of  the  tube.  The  presence  of  water  in  the  lighter 
ethereal  layer  may  be  shown  by  adding  a  small  amount  of  anhydrous 
copper  sulphate  (cf.  p.  430),  the  white  colour  of  which  changes  to  blue 
in  the  presence  of  moisture. 

In  most  cases  the  mutual  solubility  of  two  partially  miscible  liquids 
increases  with  the  temperature,  and  it  may  therefore  be  anticipated 
that  liquids  which  in  certain  proportions  form  two  layers  at  the 
ordinary  temperature  may  become  completely  miscible  at  higher 
temperatures.  This  is  the  case,  for  instance,  with  ordinary  phenol 
and  water,  which  mix  in  all  proportions  at  temperatures  above  68.4°. 

Solubility  of  Solids  in  Liquids— The  solubility  of  solids  in 
liquids  depends  upon  the  nature  of  the  solid,  the  nature  of  the  solvent, 
and  upon  the  temperature,  but  is  only  very  slightly  affected  by  change 
of  pressure.  In  all  cases  the  solubility  is  limited  but  varies  within 
very  wide  limits ;  thus  chalk  is  practically  insoluble  in  water,  but  the 
latter  solvent  can  take  up  its  own  weight  of  cane  sugar  at  the 
ordinary  temperature.  Two  general  methods  are  used  for  deter- 
mining the  solubility  of  a  solid  in  a  liquid.  According  to  the  first 
method,  the  finely  divided  solid  is  shaken  with  a  definite  volume  of 
the  liquid  at  constant  temperature  till  no  more  will  dissolve ;  part 
of  the  solution  is  then  removed  and  the  amount  of  the  dissolved 
substance  it  contains  determined.  The  second  method  depends 
upon  the  fact  that  in  the  great  majority  of  cases  the  solubility  in- 
creases as  the  temperature  is  raised.  The  solvent  is  heated  with 
excess  of  the  solid  to  a  temperature  higher  than  that  at  which  the 
solubility  is  to  be  determined,  and  is  then  cooled  to  the  required 
temperature  in  contact  with  the  solid,  when  the  excess  above  that 
required  to  form  a  saturated  solution  separates  out. 
6 


82     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


As  already  indicated,  a  saturated  solution  represents  a  state  of 
equilibrium  between  two  phases,  and  in  the  present  case  a  liquid  and 
a  solid  phase  are  in  equilibrium.  It  is  instructive  to  compare  an 
equilibrium  of  the  type  now  under  discussion  with  that  between 
water  and  water  vapour.  In  the  latter  case,  there  is  a  definite  con- 
centration of  vapour  in  equilibrium  with  water  at  a  definite  tem- 
perature, and  the  vapour  concentration  is  independent  of  the  amount 
of  vapour  (that  is  of  the  size  of  the  space  accessible  to  it)  and  of  the 
amount  of  water.  In  the  same  way,  the  concentration  of  a  saturated 
solution  in  contact  with  undissolved  solid  is  independent  of  the 
volume  of  the  solution  and  of  the  amount  of  undissolved  solid. 

The  extent  of  the  solubility  of  a  solid  in  a  liquid  under  definite 
conditions  may  be  expressed  in  a  number  of  ways  :  (i)  As  the 
number  of  parts  by  weight  (grams)  of  the  solute  present  in  (a)  100 
parts  by  weight  (grams)  or  (b}  100  parts  by  volume  (c.c.)  of  the  saturated 
solution  ;  (2)  as  the  number  of  parts  by  weight  (grams)  of  the  solute 
taken  up  by  (a)  100  parts  by  weight  (grams)  or  (b}  100  parts  by  volume 
(c.c.)  of  the  solvent.  The  second  method  appears  to  possess  certain 
advantages,  and  will  be  mainly  used  in  the  present  book. 
Effect  of  Temperature  on  the  Solubility  of  Solids  in 

Liquids— The  solubilities  of 
a  number  of  salts  in  the  same 
solvent,  water,  and  the  varia- 
tion of  the  solubilities  with 
temperature,  is  represented 
graphically  in  Fig.  32.  The 
ordinates  represent  the  num- 
ber of  grams  of  salt  taken  up 
by  100  grams  of  water  and  the 
abscissas  represent  tempera- 
tures. The  curve  for  potas- 
sium nitrate  is  very  steep, 
indicating  that  the  solubility 
of  the  salt  increases  very 
rapidly  as  the  temperature 
rises.  Thus  it  can  be  gathered 
from  the  curve  that  at  10°  100 
grams  of  water  dissolve  about 
20  grams  of  the  salt,  and  at 
50°  about  85  grams,  so  that  a 


140 
130 
120 
110 
100 
90 
'80 
70 
60 
50 
40 
30 
20 

/ 

7 

~? 

* 

0/> 

k 

1  ' 

C  - 

^ 

•  ZJ 

« 

t 

\\ 

c 

1 

^ 

? 

/ 

/ 

f 

/ 

I/ 

/ 

/ 

— 

s 

^/ 

/ 

C 

4f* 

/ 

/ 

/ 

S* 

22 

-A 

fc- 

•  •"^ 

- 

.  '      • 

i 

**} 

^ 



^ 

i  * 
:  - 

^. 

^, 
tf 

„-- 

:  = 
-  -? 

^ 

^  = 

;  =  = 

s 

r 

$ 

^ 

^  w 

?|V 

^ 

** 

3 

2 

f 

:-i 

•" 

^ 

10 
C 

— 

^  - 

=- 

:* 

»-* 

•  •  10      ZO      JO     40      W     60      70      60      90     IOC 

Temperatures . 
FIG.  32. 


rise  of  temperature  of  40°  has  more  than  quadrupled  the  solubility. 


SOLUTION 


On  the  other  hand  the  solubility  of  sodium  chloride  at  o°  is  about  36 
grams  and  at  100°  nearly  40  grams  in  100  grams  of  water,  so  that  the 
change  of  solubility  with  temperature  is  very  slight.  The  temperature 
coefficients  of  solubility  of  the  great  majority  of  salts  in  water  are 
greater  than  that  of  sodium  chloride  but  less  than  that  of  potassium 
nitrate,  as  the  diagram  indicates. 

The  solubility  of  sodium  sulphate  in  water  is  remarkable  inasmuch 
as  it  increases  fairly  rapidly  with  the  temperature  up  to  33°  and  then 
slowly  diminishes  as  the  temperature  is  further  increased  (Fig.  33). 

An  examination  of  the  solid  in  contact  with  the  solution  will  show 
that  at  below  33°  it  consists  of  a  compound  of  the  salt  and  water,  a 
hydrate  (p.  91),  above  33°  the  solid  in  equilibrium  with  the  solution 
consists  of  the  anhydrous 
salt.  We  may  therefore 
say  that  the  part  BC  of 
the  curve  represents  the 
solubility  of  a  hydrate,  the 
branch  CD  that  of  the 
anhydrous  salt. 

In  a  few  exceptional 
cases  the  solubility  of  a 
solid  in  a  liquid  diminishes 
as  the  temperature  rises. 
When  fresh  lime  is  shaken 
up  with  water  for  some 
time  at  the  ordinary  tem- 
perature and  the  excess  of 
solid  is  allowed  to  settle,  a 
clear  saturated  solution  of  calcium  hydroxide  in  water,  familiarly 
known  as  lime  water,  is  obtained.  On  warming  the  solution  it 
becomes  turbid  from  separation  of  solid,  showing  that  the  solubility 
diminishes  with  increase  of  temperature.  Some  further  examples  of 
negative  temperature-coefficients  of  solubility,  as  they  are  called,  will 
be  mentioned  in  connexion  with  the  respective  salts. 

The  effect  of  temperature  on  the  solubility  is  closely  connected 
with  the  question  whether  heat  is  given  out  or  taken  up  when  solution 
takes  place.  Certain  substances,  such  as  potassium  nitrate,  dissolve 
in  water  with  absorption  of  heat,  as  is  shown  by  the  fact  that  the 
temperature  falls  while  the  salt  is  being  dissolved ;  other  substances 
give  out  heat  during  solution  and  the  temperature  rises  (example, 
bromine  in  water).  It  has  now  been  found  that  for  substances  which 


Grams  of  salt  in  100  grams  water  -> 
6  8  8  §  3  g  s 

s 

V 

/ 

-^*. 

-•   .. 

»•   1- 

^ 

/ 

/ 

x 

? 

10        20       30       40       SO       60        70       80       90 

Temperatures  —  > 
FIG.  33. 

84     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

dissolve  with  absorption  of  heat  the  solubility  increases  as  the  tem- 
perature is  raised,  whilst  with  substances  which  dissolve  with  evolution 
of  heat  the  solubility  becomes  less  as  the  temperature  is  raised.  A 
good  illustration  of  this  rule  is  the  solubility  of  sodium  sulphate 
(Fig-  33)  5  corresponding  with  the  fact  that  the  hydrate  dissolves  with 
absorption  of  heat,  the  solubility  increases  with  rise  of  temperature 
(branch  BC),  the  anhydrous  salt,  on  the  other  hand,  dissolves  with 
evolution  of  heat,  and  its  solubility  diminishes  with  rise  of  temperature 
(branch  CD).  The  usefulness  of  this  rule  is  considefably  lessened 
by  the  fact  that  it  refers  to  the  heat  effect  of  dissolving  a  substance 
in  its  own  saturated  (or  practically  saturated)  solution  and  not  to 
whether  heat  is  given  out  or  absorbed  when  a  substance  is  dissolved 
in  the  pure  solvent.  It  sometimes  happens  that  although  heat  is 
absorbed  when  a  definite  amount  of  salt  is  dissolved  in  a  large 
volume  of  water,  heat  is  given  out  when  the  same  amount  of  salt 
is  dissolved  in  a  solution  containing  a  large  proportion  of  the  same 
salt.  The  rule  in  question  should  therefore  be  used  with  considerable 
caution  (cf.  p.  172). 

Supersaturated  Solutions— If  water  is  saturated  at  a  definite 
temperature,  say  20°,  by  shaking  with  excess  of  a  salt  the  solubility 
of  which  increases  with  rise  of  temperature,  and  a  portion  of  the 
solution,  free  from  undissolved  solid,  is  heated  to,  say,  40°,  it  will  then 
be  capable  of  dissolving  more  salt,  and  is  therefore  termed  un- 
saturated.  If,  on  the  other  hand,  a  concentrated  solution  of  the  same 
salt  is  prepared  at,  say,  40°,  and  then  allowed  to  cool  slowly  in  the 
complete  absence  of  undissolved  solid,  solid  salt  does  not  necessarily 
separate  when  the  temperature  has  fallen  below  that  at  which  the 
solution  is  saturated  in  contact  with  solid.  Such  a  solution  is  said  to 
be  supersaturated.  When  a  minute  fragment  of  the  solid  is  added  to 
such  a  solution  the  excess  of  dissolved  solid  at  once  separates,  and 
the  concentration  falls  to  that  of  a  saturated  solution  at  the  tempera- 
ture of  the  experiment. 

The  above  statements  may  be  illustrated  as  follows.  If  crystallized 
sodium  sulphate  is  heated  with  its  own  weight  of  water  till  a  perfectly 
clear  solution  is  obtained,  and  the  neck  of  the  flask  is  closed  with  a 
plug  of  cotton-wool,  the  solution  will  remain  perfectly  clear  when  the 
temperature  has  fallen  to  that  of  the  room,  although  it  is  then  highly 
supersaturated.  If,  however,  a  minute  crystal  of  the  solid  is  added, 
crystallization  at  once  takes  place. 

The  amount  of  a  solid  sufficient  to  start  crystallization  in  a  super- 
saturated solution  is  extremely  small.  According  to  Ostwald,  the 


SOLUTION  85 

almost  inconceivably  minute  amount  of  icr10  gram l  (a  ten-millionth 
part  of  a  milligram)  brings  a  supersaturated  solution  of  sodium 
sulphate  to  crystallization.  As  small  fragments  of  many  salts  are 
present  in  the  floating  dust  in  the  air  of  laboratories,  it  is  usually 
necessary  to  plug  with  cotton-wool  the  neck  of  the  flask  in  which 
a  supersaturated  solution  is  being  prepared. 

Supersaturation  is  by  no  means  confined  to  solutions  of  solids  in 
liquids.  Under  certain  conditions,  highly  supersaturated  solutions  of 
gases  in  liquids  may  be  obtained. 

It  is  evident  that  there  is  a  close  analogy  between  supersaturation 
and  the  phenomenon  of  supercooling.  Both  are  illustrations  of  the 
general  rule  that  when  the  conditions  are  favourable  for  the  appear- 
of  a  new  phase,  it  does  not  necessarily  appear  unless  a  small  amount 
of  it  is  already  present  in  the  system. 

If  a  supersaturated  solution  of  sodium  sulphate  is  cooled  sufficiently, 
the  solid  salt  begins  to  separate  without  the  necessity  of  adding  a 
small  crystal  of  the  solid  phase.  A  distinction  is  sometimes  drawn 
between  the  metastable  region,  in  which  crystallization  does  not  occur 
in  the  absence  of  a  fragment  of  the  solid  phase,  and  the  labile  region, 
in  which  crystallization  of  a  supersaturated  solution  occurs  spon- 
taneously (cf.  p.  293). 

Solution  and  the  Kinetic  Theory— The  experimental  fact 
that  a  salt  can  remain  in  equilibrium  with  a  solution  at  a  definite 
temperature,  under  which  circumstances  the  solution  is  said  to  be 
saturated,  is  instructive  when  considered  from  the  standpoint  of  the 
kinetic  theory.  We  may  suppose  that  when  a  solid  is  in  contact 
with  a  liquid  it  tends  to  send  out  particles  into  the  liquid.  These 
particles  move  about  in  the  liquid,  and  some  of  them  will  presumably 
return  and  redeposit  on  the  solid.  The  number  thus  returning  to  the 
solid  state  will  be  the  greater  the  greater  the  concentration  of  the 
dissolved  substance,  and  ultimately  a  stage  will  be  reached  when 
the  number  sent  out  is  just  balanced  by  the  number  returning  to  the 
solid— this  is  the  equilibrium  condition.  The  tendency  of  a  solid  to 
send  out  particles  into  a  liquid  in  contact  with  it  is  known  as  its 
solution  pressure ;  at  equilibrium  the  solution  pressure  is  balanced  by 
the  pressure  of  the  dissolved  substance. 

1  Very  small  amounts  are  most  conveniently  expressed  by  means  of  negative 
indices.  Thus  ro"1  gram  is  ^  gram,  io~3  gram  is  y^  gram,  and  so  on. 


CHAPTER  VIII 
CHLORINE   AND    HYDROCHLORIC   ACID 

TT)EFERENCE  has  already  been  made  to  both  these  important 
-L  V  substances  in  the  chapter  on  water.  It  has  been  pointed  out 
that  chlorine  is  an  element  which  can  be  made  to  combine  with  the 
hydrogen  of  water,  forming  hydrogen  chloride  and  setting  free 
oxygen  (p.  52).  The  chemical  equation  representing  the  change  is 
as  follows : — 

Chlorine  +  water = hydrogen  chloride  +  oxygen. 

Hydrogen  chloride  is  sometimes  called  hydrochloric  acid  gas ;  its 
solution  in  water  is  hydrochloric  acid.  For  several  reasons  it  is 
convenient  to  deal  with  these  two  substances  at  the  present  stage. 

CHLORINE 

History — Chlorine  was  discovered  in  1774  by  Scheele,  who 
obtained  it  by  heating  hydrochloric  acid  with  a  naturally  occurring 
oxide  of  manganese.  It  was  for  a  number  of  years  called  oxymuriatic 
acid,  being  supposed  to  contain  oxygen  (p.  92).  In  1810  Sir 
Humphry  Davy  proved  it  to  be  an  element,  and  gave  it  its  present 
name  in  allusion  to  its  greenish-yellow  colour  (%\  mp6s= greenish  - 
yellow). 

Occurrence — Chlorine  does  not  occur  free  in  nature,  on  account 
of  its  great  chemical  activity,  but  is  found  in  large  amount  combined 
with  sodium,  potassium,  magnesium,  and  other  metals.  Its  com- 
pound with  sodium,  known  as  sodium  chloride  or  common  salt,  is,  as 
already  mentioned,  the  salt  present  in  the  largest  proportion  in  sea- 
water,  and  is  also  found  in  mines  in  Galicia  and  elsewhere.  Potassium 
and  magnesium  chlorides  are  found  along  with  sodium  chloride  in 
salt  deposits  at  Stassfurt  in  Germany.  Both  chlorine  and  hydro- 
chloric acid  are,  however,  obtained  almost  exclusively  from  sodium 
chloride. 

Preparation — (i)  As  hydrochloric  acid  is  a  chemical  compound 
of  chlorine  and  hydrogen  which  is  easily  obtained  in  quantity,  it  is 


CHLORINE   AND   HYDROCHLORIC   ACID        87 

natural  to  use  it  as  a  source  of  chlorine.  For  this  purpose  the 
hydrogen  has  to  be  removed,  and  this  is  most  conveniently  done  by 
causing  it  to  combine  with  oxygen.  The  equation  representing  this 
action  is  as  follows — 

hydrogen  chloride  +  oxygen^twater  +  chlorine 

and  is  the  reverse  of  the  equation  on  the  previous  page,  which 
represents  the  decomposition  of  water  by  free  chlorine.  The  reversi- 
bility of  the  reaction  is  conveniently  represented  by  oppositely 
directed  arrows  (cf.  p.  164). 

The  apparatus  used  is  represented  in  Fig.  34.  A  stream  of  air  is 
passed  through  the  bottle,  which  contains  concentrated  hydrochloric 
acid,  and  the  mixture  of  air  and  hydrogen  chloride  thus  obtained  is 


Aip 


FIG.  34- 

led  through  the  bulb  tube,  which  is  strongly  heated.  The  tube 
contains  pieces  of  pumice  stone  soaked  with  a  solution  of  copper 
chloride  (or  sulphate).  The  gas  escaping  at  the  end  of  the  tube 
contains  chlorine,  which  can  be  recognized  by  its  smell  and  by  other 
characteristic  tests  mentioned  below.  It  is,  however,  mixed  with 
excess  of  air  and  of  hydrogen  chloride,  and  cannot  be  obtained  even 
approximately  pure  by  this  method. 

In  the  absence  of  copper  chloride  practically  no  chlorine  is  formed 
when  a  mixture  of  air  and  hydrogen  chloride  is  passed  through  a  red- 
hot  tube.  It  is  evident,  therefore,  that  the  copper  salt  has  accelerated 
the  change  represented  by  the  upper  arrow,  whilst  the  salt  can  be 
recovered  unchanged  at  the  end  of  the  reaction.  This  is  a  further 
illustration  of  what  we  have  already  termed  catalytic  actions. 

The  method  just  described  is  used  for  preparing  chlorine  on  the 


88     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


commercial  scale,  and  is  called  the  Deacon  process.  The  salt  of 
copper  chiefly  used  as  a  catalyst  for  the  Deacon  process  is  cuprous 
chloride  (cf.  p.  428). 

(2)  Instead  of  using  free  oxygen  to  decompose  hydrogen  chloride, 
it  is  much  more  convenient  to  use  combined  oxygen  in  the  form  of 
certain  peroxides  and  other  compounds  containing  a  large  proportion 
of  oxygen.  The  substance  generally  used  for  this  purpose  is  the 
black  powder  known  as  manganese  peroxide  or  pyrolusite,  which 

occurs  largely  in  nature.  The 
apparatus  used  for  preparing 
chlorine  by  this  method  is  repre- 
sented in  Fig.  35.  The  mixture 
of  pyrolusite  and  a  concentrated 
aqueous  solution  of  hydrochloric 
acid  is  heated  gently  in  a  large 
flask  A,  and  the  issuing  chlorine 
is  first  passed  through  water  in  a 
two-necked  bottle  B  (the  so-called 
Woulf's  bottle)  to  free  it  from 
hydrogen  chloride,  and  then,  if 
required  perfectly  dry,  through  a 
second  bottle  containing  concen- 
trated sulphuric  acid.  As  chlorine 
is  fairly  soluble  in  water,  and  acts 
chemically  on  mercury,  it  cannot 
be  collected  over  either  of  these 
liquids.  As  it  is  considerably 
heavier  than  air,  it  can  readily  be 
collected  by  allowing  it  to  issue 
from  the  delivery  tube  at  the 
bottom  of  a  gas  jar  C,  as  shown  in 


FIG.  35. 


the  figure.     This  method  of  collecting  heavy  gases  by  upward  dis- 
placement of  air  has  already  been  mentioned  (p.  39). 

The  equation  representing  the  chemical  change  just  described  is  as 
follows  : — 

Manganese  peroxide + hydrochloric  acid 
=  chlorine  +  manganese  chloride  +  water. 

The  manganese  chloride  which,  as  its  name  indicates,  is  a  chemical 
compound  of  manganese  and  chlorine,  remains  behind  in  the  flask. 
A  modification  of  this  method,  due  to  Weldon  (p.  545),  may  be  used 
for  the  commercial  preparation  of  chlorine. 


CHLORINE   AND    HYDROCHLORIC   ACID        89 

(3)  Instead  of  manganese  peroxide,  chlorinated  lime,  lead  peroxide, 
potassium   bichromate  or  potassium   permanganate,  all   compounds 
containing    oxygen,   may   be    used   to   effect   the   decomposition  of 
hydrochloric  acid.      The  reactions  are  described  in  connexion  with 
the    compounds    themselves.      Chlorinated    lime,   otherwise    called 
bleaching  powder,  is  particularly  suitable,  as  the  chlorine  is  rapidly 
given  off  in  the  cold. 

(4)  By  electrolysis.     Chlorine  can  also  be   prepared  from  hydro- 
chloric acid  by  electrolysis  in  concentrated  aqueous  solution.      An 
apparatus  similar  to  that  shown  on  p.  14  may  be  used  ;  owing,  how- 
ever, to  the  fact  that  chlorine  attacks  platinum,  carbon  electrodes  must 
be  used.     The  gas  which  comes  off  at  the  positive  pole  is  not  pure 
chlorine,  but  always  contains  oxygen,  and  the  proportion  of  the  latter 
gas  is  the  greater  the  more  dilute  the  hydrochloric  acid  solution. 
The  oxygen  probably  results  from  the  action  of  the  freshly  liberated 
chlorine  on  the  water. 

Compounds  of  chlorine  with  metals,  the  so-called  metallic  chlorides, 
also  yield  chlorine  on  electrolysis.  On  the  commercial  scale,  the 
preparation  of  chlorine  by  electrolysis,  e.g.  of  sodium  chloride  in 
aqueous  solution  (p.  398),  has  almost  entirely  displaced  the  chemical 
methods  formerly  in  use. 

Physical  Properties — Chlorine  is  a  yellowish-green  gas  with 
a  powerful,  disagreeable  odour.  It  attacks  the  mucous  membranes 
strongly,  and  if  inhaled  in  considerable  amount  causes  death  by 
suffocation.  Its  density,  referred  to  air  as  unity,  is  about  2.45, 
referred  to  hydrogen  about  35.5.  One  litre  of  chlorine,  at  normal 
temperature  and  pressure,  weighs  3.22  grams.  Chlorine  is  very  easily 
liquefied,  at  o°  six  atmospheres  pressure  is  required,  and  at  —34° 
one  atmosphere  pressure  is  sufficient.  The  last  statement  indicates 
that  the  boiling-point  of  liquid  chlorine  is  -  34°.  At  -  102°  it  solidifies  ; 
both  the  liquid  and  the  solid  have  a  yellow  colour.  Liquid  chlorine 
is  now  a  commercial  article,  being  conveyed  in  steel  cylinders  lined 
inside  with  lead.  The  critical  temperature  of  chlorine  is  146°,  and  its 
critical  pressure  94  atmospheres. 

Chlorine  is  fairly  soluble  in  water  ;  at  room  temperature  (18°)  I 
volume  of  water  takes  up  about  2.2  volumes  of  the  gas.  The  variation 
of  the  absorption  coefficient,  a,  with  temperature,  is  represented  by 

the  formula — 

( 

«= 3.0361  -0.046 1 96/4-  o.ooo  1 1 07/2. 
The  aqueous  solution  has  the  odour  and  colour  of  the  gas,  and  is 


9o      A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

known  as  chlorine  water.  It  is  very  unstable,  especially  in  sunlight, 
hydrogen  chloride  and  oxygen  being  ultimately  formed  (see  below). 
This  may  be  shown  by  rilling  a  small  retort  with  chlorine  water  and 
exposing  it  to  bright  sunlight  while  supporting  it  in  the  position  shown 
(Fig.  36).  In  a  short  time  a  gas,  which  proves  to  be  oxygen,  collects 
in  the  upper  part  of  the  retort. 

Chemical  Properties— Chlorine  is  characterized  by  great 
chemical  activity,  combining  at  room  temperature  with  a  number 
of  other  elements  with  evolution  of  light  and  heat.  Finely  powdered 
arsenic  and  antimony,  when  shaken  into  the  gas,  at  once  catch  fire, 
forming  white  fumes  of  the  respective  chlorides.  Finely  divided 
copper,  lead,  tin,  and  "  Dutch  metal,"  which  consists  of  a  mixture  of 
copper  and  zinc  hammered  out  into  thin  sheets  like  gold  leaf,  also  catch 
fire  in  chlorine,  and  phosphorus  at  first  melts  and  then  burns  feebly 
in  the  gas.  Chlorine  does  not  unite  directly  with  carbon,  with  nitrogen, 
or  with  oxygen. 

Reference  has  already  been  made  to  the  great  affinity  between 
chlorine  and  hydrogen.  A  jet  of  hydrogen 
burning  in  the  air  will  continue  to  burn 
when  lowered  into  a  jar  of  chlorine,  and 
in  the  same  way  a  gas  jet  or  lighted  candle 
continues  to  burn  in  chlorine,  hydrogen 
chloride  being  formed,  and  the  carbon  set 
free  in  the  form  of  soot.  The  same  facts 

may  be  illustrated  very  strikingly  by  soaking  a  strip  of  filter-paper 
with  turpentine,  which  is  a  chemical  compound  of  carbon  and  hydro- 
gen, and  suspending  it  in  a  jar  of  chlorine.  In  a  few  moments  the 
turpentine  catches  fire  and  dense  clouds  of  soot  are  formed. 

It  is  a  remarkable  fact,  which  has  not  yet  been  adequately  ex- 
plained, that  many  of  the  reactions  just  described  do  not  take  place 
if  the  chlorine  is  perfectly  dry.  Thus  Dutch  metal  is  not  affected 
when  put  into  a  jar  of  chlorine  which  has  been  dried  by  means  of 
concen  trated  sulphuric  acid.  We  shall  meet  later  with  many  examples 
of  the  influence  of  traces  of  moisture  on  chemical  reactions.  The 
effects  belong  to  the  class  of  catalytic  phenomena. 

Bleaching  Action  of  Chlorine— If  chlorine  gas  is  passed 
into  a  solution  of  a  colouring  matter,  such  as  indigo  or  litmus,  the 
colour  rapidly  disappears  and  is  said  to  be  bleached.  Further,  if  a 
piece  of  Turkey  red  cloth  is  damped  and  put  into  a  jar  of  chlorine  it 
is  rapidly  bleached,  but  if  it  is  previously  carefully  dried  by  keeping 
it  for  some  time  in  a  vessel  over  concentrated  sulphuric  acid,  and  is 


CHLORINE   AND    HYDROCHLORIC   ACID        91 

then  suspended  in  a  jar  of  chlorine  gas  which  has  stood  for  an  hour 
or  two  in  a  carefully  closed  gas  jar  containing  a  layer  of  sulphuric 
acid  at  the  bottom,  no  bleaching  occurs.  As  the  presence  of  moisture 
is  thus  shown  to  be  essential  for  bleaching  to  occur,  it  is  sometimes 
assumed  that  the  actual  bleaching  effect,  the  conversion  of  a  coloured 
to  a  colourless  substance,  is  due  to  the  action  of  "nascent"  (freshly 
liberated)  oxygen  set  free  from  water  by  the  action  of  chlorine.  It 
appears,  however,  that  when  chlorine  is  dissolved  in  water  at  the 
ordinary  temperature  an  equilibrium  is  set  up  represented  by  the 
equation  (cf.  p.  179) 

Chlorine +water^hydrochloric  acid + hypochlorous  acid. 

The  hypochlorous  acid  readily  gives  up  oxygen  to  oxidizable  sub- 
stances, and  thus  acts  as  a  powerful  oxidizing  agent ;  in  sunlight  it 
decomposes  rapidly  into  hydrochloric  acid  and  free  oxygen.  The 
oxidizing  effect  of  chlorine  in  the  presence  of  moisture  is  thus 
satisfactorily  explained. 

Chlorine  bleaches  ordinary  ink,  but  has  little  or  no  action  on 
printer's  ink.  The  colour  of  the  latter  is  due  to  particles  of  carbon, 
which  at  the  ordinaiy  temperature  is  not  acted  on  either  by  oxygen 
or  chlorine. 

Chlorine  has  also  a  powerful  destructive  action  on  low  organisms, 
such  as  disease  germs,  and  is  therefore  largely  used  as  a  disinfectant. 
This  effect  is  probably  also  connected  with  the  intermediate  formation 
of  hypochlorous  acid. 

Chlorine  is  also  used  commercially  in  the  extraction  of  gold  from 
its  ores,  as  it  forms  a  readily  soluble  chloride  with  this  metal  (p.  428), 

Chlorine  Hydrate — When  chlorine  gas  is  passed  foj-  some  time 
into  water  cooled  to  o°,  a  greenish  crystalline  substance,  consisting  of 
chlorine  and  water  in  chemical  combination,  separates.  The  crystal- 
line substance  is  called  chlorine  hydrate.  It  is  fairly  stable  at  low 
temperatures,  but  above  9.6°  splits  up  into  water  and  chlorine.  The 
composition  of  chlorine  hydrate  will  be  referred  to  later  (p.  112). 
Faraday  first  obtained  liquid  chlorine  by  heating  chlorine  hydrate  in 
one  end  of  the  tube  represented  in  Fig.  28  and  placing  the  other  end 
in  a  freezing  mixture  (cf.  p.  71). 

The  name  hydrate  is  applied  to  a  substance  formed  by  the  com- 
bination of  water  as  a  whole  with  other  compounds,  or,  as  in  the 
present  case,  with  elements.  An  enormous  number  of  well-known 
compounds,  such  as  common  salt,  copper  sulphate,  sodium  carbonate, 
etc.,  form  hydrates  with  water,  but  very  few  hydrates  are  known  which 
have  an  element  as  one  of  the  components. 


92     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

HYDROCHLORIC  ACID  (HYDROGEN  CHLORIDE) 

History — The  aqueous  solution  of  hydrogen  chloride  known  as 
hydrochloric  acid  has  been  known  since  the  middle  of  the  seventeenth 
century.  It  was  obtained  by  heating  sodium  chloride  (common  salt) 
with  sulphuric  acid,  the  method  which  is  still  employed  on  the  com- 
mercial scale.  It  was  formerly  known  as  spirit  of  salt^  and  also  as 
muriatic  acid.  As  chlorine  is  usually  prepared  by  oxidizing 
hydrochloric  acid,  and  was  formerly  supposed  to  contain 
oxygen,  it  is  now  easy  to  see  why  it  used  to  be  called  oxy- 
muriatic  acid  (p.  86).  The  mistake  was  at  the  time  a  very 
natural  one,  but  we  know  now  that  oxidizing  agents  do  not  add 
„  on  oxygen  to  hydrogen  chloride,  but  remove  hydrogen  from  it. 
The  gas  itself,  hydrogen  chloride,  was  discovered  by  Priestley 
(1774),  who  was  the  first  to  collect  gases  over  mercury. 

Preparation — (i)  Hydrogen  chloride  can  be  prepared  by 
direct  combination  of  its  elements,  or  by  the  action  of  chlorine 
on    many    compounds    containing   combined    hydrogen,    as 
A>iJL    already  mentioned. 

VjT        The  combination  of  the  elements  can  be  shown  very  instruc- 
tively in  a  thick-walled  tube  (Fig.  37)  provided  with  a  stopcock 
at  each  end,  and  divided  into  two  equal  parts  by  a  stopcock 
in  the  middle.     The  two  halves  of  the  tube  are  filled  in  red 
light  with   dry  hydrogen  and  chlorine  respectively  by  dis- 
placement of  mercury,  the  central  stopcock  is  opened  to  allow 
the  gases  to  mix,  and  the  tube  is  then  exposed  to  diffused 
daylight.     The  greenish  colour  of  the  chlorine  gradually  dis-4 
appears,  and  after  a  time  it  will  be  found  that  the  gases  have 
combined  completely  to  form  hydrogen  chloride.     If  then  one 
of  the  end  stopcocks  is  opened  carefully  under  mercury,  no 
gas  will  escape  and  no  mercury  will  enter.     This  experiment 
'  37' illustrates    two    very   important  points:    (i)   hydrogen   and 
chlorine   combine    in   equal  volumes  to  form   hydrogen  chloride ; 
(2)  hydrogen  and  chlorine   combine   without  change  of  volume  to 
form  hydrogen  chloride. 

If  equal  volumes  of  the  mixed  gases  are  exposed  to  bright  sunlight 
or  to  a  magnesium  flash-light,  combination  takes  place  explosively. 
This  experiment  should  not  be  tried  in  the  tube  just  described,  but 
may  be  done  quite  safely  by  supporting  a  sealed  thin-walled  bulb, 
filled  with  the  mixed  gases,  behind  a  glass  screen  and  exposing  it  to 
a  magnesium  flash-light.  Combination  is  practically  instantaneous, 


CHLORINE   AND   HYDROCHLORIC   ACID        93 

and  the  bulb  is  reduced  to  fragments.  The  preparation  of  the  bulbs 
is  effected  by  blowing  a  few  of  them  from  a  single  piece  of  glass  tube 
in  such  a  way  that  they  remain  connected  by  thin  capillary  tubes.  A 
stream  of  mixed  hydrogen  and  chlorine,  prepared  by  electrolysis,  is 
passed  through  the  series  of  bulbs  for  a  considerable  time  in  dark- 
ness, and  then  the  connexions  between  them  are  sealed  off  with  a 
small  flame.  They  can  also  be  bought  filled  and  ready  for  use. 

All  kinds  of  light  are  not  equally  effective  in  causing  combination 
of  hydrogen  and  chlorine.  Blue  and  violet  rays  are  most  active  in 
this  respect,  whereas,  as  already  indicated,  red  light  is  practically 
without  effect.  The  combination  of  the  gases  can  also  be  brought 
about  by  passing  an  electric  spark  or  by  applying  a  lighted  taper  to 
the  mouth  of  a  jar  containing  them,  as  in  the  case  of  a  mixture  of 
hydrogen  and  oxygen. 

The  fact  that  hydrogen  and  chlorine  combine  slowly  under  the  in- 
fluence of  light  is  taken  advantage  of  in  the  construction  of  one  form 
of  actinometer^  an  instrument  for  measuring  the  chemical  activity  of 
light.  The  gases  are  confined  over  water  and  kept  at  constant 
pressure,  and  from  the  rate  of  diminution  of  volume  (the  hydrogen 
chloride  which  is  formed  at  once  dissolves  in  the  water)  the  chemical 
activity  of  a  source  of  light  can  be  estimated. 

It  is  a  curious  fact  that  when  the  gases  are  exposed  to  light, 
practically  no  combination  takes  place  for  a  short  time ;  after- 
wards the  rate  of  combination  is  quite  regular.  This  initial  period 
of  no  reaction  is  termed  the  "  period  of  induction."  In  spite  of  a 
great  deal  of  work  on  the  subject,  the  phenomenon  is  by  no  means 
understood. 

(2)  Hydrogen  chloride  is  almost  invariably  prepared,  both  in  the 
laboratory  and  on  the  commercial  scale,  by  the  action  of  sulphuric 
acid  on  common  salt.  When  the  acid  is  not  used  in  too  great  excess, 
the  chemical  change  is  represented  by  the  equation 

sodium  chloride  +  sulphuric  acid  =  sodium  sulphate  4- 
hydrogen  chloride. 

The  apparatus  used  is  similar  to  that  represented  in  Fig.  35.  Dry 
sodium  chloride  is  placed  in  a  flask  provided  with  a  thistle  funnel,  and 
concentrated  sulphuric  acid  is  added  through  the  funnel.  On  gently 
warming,  hydrogen  chloride  is  given  off  in  a  steady  stream  ;  it  can  be 
dried  by  passing  through  a  bottle  containing  pumice  stone  soaked  with 
concentrated  sulphuric  acid,  and  then  collected  over  mercury  or  by 
upward  displacement  of  air  (p.  39).  As  the  gas  escapes  into  the  air 


94     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

white  fumes  will  be  noticed  ;  these  are  due  to  the  combination  of  the 
gas  with  the  moisture  of  the  air. 

As  already  mentioned,  hydrogen  chloride  is  usually  employed  in  the 
form  of  its  solution  in  water,  known  as  hydrochloric  acid.  A  method 
of  preparing  such  a  solution  is  shown  in  Fig.  44.  The  glass  tube 
dips  just  below  the  surface  of  the  water,  and  in  order  to  prevent  the 
water  from  being  .sucked  back  into  the  flask,  the  glass  tube  has  a  bulb 
in  the  middle.  In  such  a  case,  when  the  entrance  of  water  into  the 
generating  flask  might  cause  an  accident,  it  is  well  to  interpose  an 
empty  bottle  between  the  flask  and  the  absorption  vessel. 

Hydrogen  chloride  may  be  obtained  very  conveniently  for  labora- 
tory purposes  by  allowing  concentrated  sulphuric  acid  to  drop  from 
a  funnel  provided  with  a  stopcock  into  commercial  concentrated 
hydrochloric  acid.  An  apparatus  similar  to  that  represented  in 
Fig.  44  may  be  used.  Little  or  no  heating  is  required. 

The  commercial  preparation  of  hydrochloric  acid  is  again  referred 
to  in  connexion  with  sodium  carbonate  (p.  404). 

Physical  Properties — Hydrogen  chloride  is  a  colourless  gas 
with  a  suffocating  odour.  Its  density  referred  to  hydrogen  is  18.3,  so 
that  it  is  about  1.26  times  heavier  than  air.  By  cooling  under  pressure 
it  can  readily  be  converted  into  a  colourless  liquid  which  boils  at 
—  83° ;  in  other  words  the  vapour  pressure  of  liquid  hydrogen  chloride 
at  —  83°  is  one  atmosphere.  At  o°,  the  vapour  pressure  is  equal  to  26 
atmospheres.  Its  critical  temperature  is  52°,  so  that  it  can  be  liquefied 
by  pressure  alone  at  the  ordinary  temperature ;  its  critical  pressure 
is  86  atmospheres. 

One  of  the  most  remarkable  properties  of  hydrogen  chloride  is  its 
great  solubility  in  water.  This  may  be  illustrated  in  a  striking  way 
by  passing  a  few  drops  of  water,  by  means  ot  a  bent  pipette,  into  a 
quantity  of  the  gas  confined  over  mercury.  The  mercury  will  be 
observed  to  rise  rapidly  in  the  tube,  showing  that  water  can  absorb 
many  times  its  own  volume  of  the  gas.  A  still  more  striking  method 
of  illustrating  the  same  fact  is  indicated  in  Fig.  38.  The  flask  is 
completely  filled  with  hydrogen  chloride  through  the  long  glass  tube, 
which  is  open  at  both  ends  ;  the  outer  end  of  the  tube  is  then  placed 
in  a  vessel  of  water,  as  shown,  and  the  liquid  sucked  up  the  central 
tube  by  means  of  the  side  tube.  When  the  water  just  reaches  the 
top  of  the  inner  tube  the  side  tube  is  closed.  The  first  few  drops 
of  water  dissolve  nearly  all  the  gas  and  produce  a  partial  vacuum, 
into  which  the  water  is  forced  by  the  pressure  of  the  atmosphere. 

At  o°  i  volume  of  water  absorbs  503  volumes  of  hydrogen  chloride^ 


CHLORINE   AND   HYDROCHLORIC   ACID        95 


but  the  solubility  diminishes  rapidly  with  rise  of  temperature.  The 
effect  of  pressure  on  the  solubility  of  hydrogen  chloride  in  water  is 
not  represented  by  Henry's  law ;  as  a  matter  of  fact,  the  amount 
dissolved  increases  only  slightly  with  considerable  increase  of 
pressure.  This  is  illustrated  by  the  following  table,  in  which  the 
upper  line  represents  the  pressure  in  mms.  of  mercury  and  the 
lower  line  the  weight  of  acid  dissolved  by  100  grams  of  water  at 
o°  (Roscoe). 


Grams   of  hydro-  ) 
gen  chloride.       ) 


100 

65.7 


200 
70.7 


300 

73-8 


500 

78.2 


700 
81.7 


IOOO 


85.6 


It  has  already  been  stated  (p.  79)  that  Henry's  law  does  not  hold  for 
very  soluble  gases. 

At  15°  water  takes  up  about  43  per  cent,  by 
weight  of  hydrogen  chloride  under  a  pressure 
of  the  gas  equal  to  760  mm.  of  mercury,  and 
this  is,  of  course,  the  most  concentrated  acid 
which  can  be  obtained  under  ordinary  con- 
ditions ;  its  density  is  1.21.  The  concentrated 
acid  of  commerce  contains  about  36.5  per  cent, 
by  weight  of  the  acid,  the  density  of  the  solu- 
tion is  i. i 6. 

When  the  concentrated  aqueous  solution  is 
heated,  hydrogen  chloride  alone  is  given  off  at 
first,  and  the  concentration  of  the  solution 
gradually  diminishes  till  it  reaches  20.2  per  cent, 
when  the  remainder  of  the  mixture  distils  un- 
changed in  composition  at  110°.  If,  on  the 


FIG.  38. 


other  hand,  a  dilute  solution  is  boiled,  only  water  is  given  off  at  first, 
and  the  concentration  of  the  solution  in  the  distilling  flask  gradually 
increases  (with  simultaneous  rise  of  temperature)  till  it  again  reaches 
20.2  per  cent.,  when  the  residue  distils  at  constant  temperature 
as  before.  If  we  happen  to  start  with  a  20.2  per  cent,  solution 
the  whole  of  it  distils  at  constant  temperature,  like  a  pure  liquid 
(p.  64),  and  the  distillate  is  throughout  of  the  same  composition  as 
the  solution  in  the  distilling  flask. 

The  behaviour  of  mixtures  of  hydrogen  chloride  and  water  on  dis- 
tillation is  different  from  that  of  the  great  majority  of  binary  mixtures 
of  liquids.  We  have  already  seen  that  a  liquid  boils  when  its  vapour 
pressure  is  equal  to  that  of  the  atmosphere,  and  in  the  same  way 


96     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

a  binary  mixture  boils  when  the  sum  of  the  partial  pressures  of  its 
components  is  equal  to  atmospheric  pressure.  In  a  mixture  of  two 
liquids  of  different  boiling-points,  one  of  the  components  has  in  the 
great  majority  of  cases,  not  only  a  lower  vapour  pressure  than  the 
other,  but  a  lower  vapour  pressure  than  a  mixture  of  the  components 
in  any  proportions.  When  such  a  mixture  is  heated,  the  less  volatile 
component  will  tend  to  remain  behind,  and  a  partial  separation  can 
thus  be  effected,  and  this  will  take  place  the  more  readily  the  greater 
the  difference  in  the  boiling-points  of  the  components.  If,  however, 
a  mixture  of  the  components  in  a  certain  definite  proportion  happens 
to  have  a  lower  vapour  pressure  than  that  of  either  of  the  components 
or  of  a  mixture  of  the  two  in  any  other  proportion,  it  is  evident  that 
on  distillation  a  mixture  of  this  composition  will  tend  to  remain  behind. 
We  have  an  excellent  illustration  of  the  last  case  in  the  mixture 
of  hydrogen  chloride  and  water  containing  20.2  per  cent,  of  the 
acid,  and  the  reason  why  this  mixture  tends  to  remain  behind  on 
distillation,  and  finally  distils  at  constant  temperature,  will  now  be 
understood. 

It  was  long  thought  that  this  constant-boiling  mixture  of  hydrogen 
chloride  and  water  is  a  definite  chemical  compound,  but  Roscoe  and 
Dittmar  (1860)  showed  that  the  composition  of  the  mixture  depends 
on  the  pressure.  Thus  at  a  pressure  of  250  cm.  of  mercury  the 
constant-boiling  mixture  contains  18  per  cent.,  at  76  cm.  (atmos- 
pheric pressure)  20.2  per  cent,  (boiling-point  no0),  and  at  10  cm. 
22.9  per  cent,  of  hydrogen  chloride  (boiling-point  62°). 

Chemical  Properties— The  most  striking  feature  with  refer- 
ence to  the  chemical  properties  of  hydrogen  chloride  is  the  remark- 
able difference  in  activity  between  the  liquefied  gas  and  the  aqueous 
solution.  Whilst  the  solution  of  the  gas  in  water  turns  KUnus  red,  at 
once  acts  on  lime,  and  dissolves  such  metals  as  zinc  and  magnesium, 
the  liquefied  gas,  in  the  entire  absence  of  moisture,  shows  none  of 
these  properties.  It  is  evident,  therefore,  that  only  the  aqueous 
solution  acts  as  an  acid  ;  anhydrous  hydrogen  chloride  has  no  acid 
properties.  The  explanation  of  this  remarkable  alteration  of  pro- 
perties in  contact  with  water  will  be  given  later. 

The  aqueous  solution  is  a  typical  acid  ;  it  turns  red  litmus  blue, 
has  a  sour  taste,  and,  like  sulphuric  acid  (p.  35),  acts  on  metals  such 
as  magnesium  and  zinc,  hydrogen  being  set  free.  The  general 
properties  of  acids  are  further  considered  below  (p.  98). 

Composition  of  Hydrogen  Chloride— The  composition  by 
volume  of  hydrogen  chloride,  like  that  of  water,  may  be  determined 


CHLORINE   AND   HYDROCHLORIC   ACID        97 

by  analytical  or  by  synthetical  methods.  It  has  already  been  shown 
by  a  synthetical  method  (the  double  tube  experiment,  p.  92)  that  the 
hydrogen  and  chlorine  combine  in  equal  volumes  to  form  hydrogen 
chloride,  and  that  by  the  combination  of  one  volume  of  hydrogen  and 
one  volume  of  chlorine  two  volumes  of  hydrogen  chloride  are  formed. 
That  combination  is  complete  when  equal  volumes  of  the  gases  are 
taken  can  be  shown  by  cautiously  opening  the  lower  stopcock  under 
water,  when  the  latter  will  enter  and  completely  fill  the  tube. 

The  conclusions  drawn  from  this  experiment  can  be  confirmed  by 
an  analytical  method  described  by  Roscoe.  The  apparatus  used  is 
represented  in  Fig.  39,  and  consists  essentially  of  a 
(J-tube,  one  limb  of  which  is  open,  and  the  other  is 
provided  with  a  stopcock.  Both  limbs  are  at  first 
filled  with  mercury,  and  then  dry  hydrogen  chloride 
is  drawn  into  the  left-hand  limb  until  it  is  about  two- 
thirds  full  at  atmospheric  pressure.  Some  liquid 
sodium  amalgam  (a  solution  of  sodium  in  mercury, 
p.  397)  is  then  poured  into  the  open  limb  of  the 
(J-tube,  the  end  of  the  latter  firmly  closed  with  the 
thumb,  the  hydrogen  chloride  transferred  to  the 
right-hand  limb  by  inclining  the  tube  and  kept  for 
a  little  time  in  contact  with  the  amalgam.  Under 
these  circumstances  the  hydrogen  chloride  is  com- 
pletely decomposed,  the  chlorine  combining  with  the 
sodium  to  form  sodium  chloride,  and  the  hydrogen 
is  set  free.  The  latter  is  transferred  to  the  left-hand 
limb,  the  pressure  again  adjusted  to  that  of  the 
atmosphere,  and  it  will  then  be  found  that  the 
hydrogen  occupies  exactly  half  the  volume  of  the 
original  hydrogen  chloride.  The  residual  gas  may  be  shown  to  be 
hydrogen  by  driving  it  out  at  the  stopcock  and  igniting. 

It  still  remains  to  find  the  volume  of  chlorine  which  has  disappeared. 
This  is  done  by  taking  a  long  narrow  glass  tube  provided  with  a  stop- 
cock at  each  end,  filling  it  with  the  mixed  gases  obtained  by  electro- 
lysis of  concentrated  hydrochloric  acid  (p.  89)  and  closing  both  stop- 
cocks. One  end  is  then  dipped  into  a  solution  of  potassium  iodide 
and  the  stopcock  opened,  when  it  will  be  found  that  the  solution  enters 
the  tube  and  occupies  exactly  half  the  volume.  Free  chlorine  acts 
upon  potassium  iodide,  forming  potassium  chloride  and  setting  free 
iodine,  which  dissolves  in  the  solution,  producing  a  brown  colour. 
The  residual  gas  is  hydrogen.  It  follows  that  the  chlorine  which  has 
7 


98     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

disappeared  occupied  a  volume  equal  to  that  of  the  hydrogen,  and 
by  combining  this  result  with  that  of  the  previous  experiment  it  is 
evident  (i)  that  hydrogen  chloride  is  entirely  made  up  of  equal 
volumes  of  hydrogen  and  chlorine  ;  (2)  that  the  gases  combine  without 
change  of  volume  to  form  hydrogen  chloride. 

From  the  densities  and  combining  volumes  of  the  gases  it  can  be 
calculated  that  hydrogen  chloride  contains  35.46  parts  of  chlorine 
to  1.0078  parts  of  hydrogen  by  weight.  Direct  determinations  of  the 
composition  by  weight  lead  to  the  same  result.  We  have  here  a 
further  illustration  of  the  law  that  definite  chemical  compounds  are 
of  constant  composition. 

The  composition  of  water  and  of  hydrogen  chloride  have  been 
considered  in  some  detail  on  account  of  their  fundamental  importance 
for  the  theory  of  chemistry. 

Acids,  Bases  and  Salts— Reference  has  already  been  made 
on  several  occasions  to  acids,  and  under  this  heading  we  have  grouped 
substances  which  have  certain  properties  in  common.  The  more 
important  of  these  properties  are  (i)  the  power  of  turning  litmus  red ; 
(2)  a  sour  taste  ;  (3)  the  liberation  of  hydrogen  when  they  are 
brought  into  contact  with  certain  metals,  such  as  zinc  or  magnesium. 
It  follows  from  the  statement  of  the  third  property  that,  as  already 
pointed  out,  all  acids  contain  hydrogen. 

It  must  be  carefully  remembered,  however,  that  a  substance  con- 
taining hydrogen  is  not  necessarily  an  acid.  Thus  liquefied  hydrogen 
chloride,  although  it  contains  hydrogen,  has  no  acid  properties,  nor 
has  water  itself.  The  full  discussion  of  this  important  subject  is 
postponed  to  a  later  stage. 

We  have  also  learned  that  when  metallic  sodium  is  added  to  water, 
a  solution  is  obtained  which  has  a  soapy  feel  and  turns  red  litmus 
blue.  These  properties  indicate  the  presence  of  a  base  in  the  solution. 
The  particular  base  in  this  case  is  sodium  hydroxide  ;  it  contains 
sodium,  hydrogen,  and  oxygen,  and  can  be  obtained  as  a  white 
solid  on  evaporating  the  solution.  A  solution  which  turns  litmus 
blue  is  said  to  have  an  alkaline  reaction.  A  solution  which  turns 
litmus  red  has  an  acid  reaction. 

If  now  to  a  solution  of  sodium  hydroxide,  containing  a  little  litmus, 
a  dilute  solution  of  hydrochloric  acid  is  added  a  little  at  a  time,  the 
mixture  being  stirred  after  each  addition,  the  blue  colour  will  remain 
for  some  time,  but  after  a  definite  amount  of  hydrochloric  acid  has 
been  added,  it  will  suddenly  turn  red.  Experiment  will  show  that 
a  drop  or  two  of  acid  or  alkali  is  sufficient  to  change  the  colour  from 


CHLORINE   AND    HYDROCHLORIC   ACID        99 

blue  to  red  or  vice  versa,  but  a  point  can  be  reached  at  which  the 
solution  is  neither  definitely  blue  or  red,  but  violet.  Such  a  solution 
has  none  of  the  properties  of  an  acid  or  a  base,  it  is  said  to  be  neutral. 
On  evaporating  off  the  water,  crystals  of  sodium  chloride  (common 
salt)  are  obtained. 

The  formation  of  a  substance  with  new  properties  is,  of  course, 
evidence  of  a  chemical  action.  We  therefore  obtain  the  very  im- 
portant result  that  when  a  typical  acid,  such  as  hydrochloric  acid,  is 
brought  in  contact  with  a  typical  base,  such  as  sodium  hydroxide,  the 
chemical  change  results  in  the  complete  disappearance  of  the  acid 
and  basic  properties,  and  a  new  substance,  with  entirely  different 
properties,  results.  The  process  just  considered  is  termed  neutraliza- 
tion, and  a  substance  resulting  from  the  neutralization  of  an  acid  by 
a  base  is  termed  a  salt.  The  statements  just  made  are  quite  general ; 
when  any  acid  is  brought  in  contact  with  any  base  a  salt  is  formed. 
In  the  present  example  we  have  seen  that  sodium  chloride  is  one  of 
the  products  of  the  neutralization.  As  it  contains  only  sodium  and 
chlorine,  there  remains  the  hydrogen  of  the  hydrochloric  acid  and  the 
oxygen  and  hydrogen  of  the  sodium  hydroxide  to  be  accounted  for. 
These  are,  however,  the  components  of  water,  and,  as  a  matter  of 
fact,  water  is  the  second  product  of  the  reaction.  The  complete 
equation  is  therefore — 

Hydrochloric  acid  +  sodium  hydroxide  =  sodium  chloride  4-  water. 

When  an  acid  is  neutralized  by  a  base,  water  is  always  one  of  the 
products. 

The  formation  of  the  salt  may  be  regarded  from  a  rather  different 
point  of  view,  as  resulting  from  the  displacement  of  the  hydrogen  of 
the  acid  by  a  metal,  in  this  case  sodium.  It  should,  however,  be 
mentioned  that  if  the  hydrogen  of  an  acid  is  only  partially  replaced 
by  a  metal,  the  resulting  substance  is  nevertheless  termed  a  salt.  A 
salt  may,  therefore,  be  defined  as  a  substance  formed  by  the  complete  or 
partial  displacement  of  the  hydrogen  of  an  acid  by  a  metal.  It  does 
not,  of  course,  matter  whether  the  displacement  is  effected  directly  by 
the  metal  itself  or,  as  in  the  present  case,  by  a  compound  of  the 
metal.  We  have  already  met  with  a  number  of  instances  of  the 
direct  displacement  of  hydrogen  from  an  acid  by  a  metal  with 
formation  of  a  salt. 

It  has  already  been  mentioned  that  certain  oxides  form  acids  with 
water,  and  are  therefore  known  as  acidic  oxides  or  anhydrides.  Other 
oxides,  of  which  calcium  oxide  (lime)  is  a  type,  when  shaken  up  with 


ioo     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

water,  yield  solutions  which  have  all  the  properties  of  a  base,  and  for 
this  reason  they  are  termed  basic  oxides.  Basic  oxides,  like  bases, 
form  salts  with  acids,  and  water  is  the  other  product  of  the  action. 
The  equation  in  the  case  of  calcium  oxide  and  hydrochloric  acid  is  as 
follows  : — 

Calcium  oxide  +  hydrochloric  acid  =  calcium  chloride  +  water. 

Salts  may  therefore  be  formed  by  the  action  of  acids  on  metals,  bases, 
or  basic  oxides. 

It  has  been  pointed  out  that  sodium  chloride  in  aqueous  solution  is 
neutral  to  litmus.  The  same  is  true  of  many  other  salts,  such  as 
calcium  chloride  and  sodium  sulphate,  but  is  by  no  means  generally 
true.  The  aqueous  solutions  of  some  salts  have  an  acid  reaction, 
those  of  others  an  alkaline  reaction. 

Oxides  of  Chlorine — Although  chlorine  cannot  be  made  to 
combine  directly  with  free  oxygen,  yet  two  oxides  of  chlorine  can  be 
obtained  by  indirect  methods.  These  oxides  are  briefly  referred  to 
here  in  connexion  with  the  fundamental  laws  of  chemistry  discussed 
in  the  next  chapter,  and  will  be  dealt  with  more  fully  in  Chapter  XIV. 
in  connexion  with  other  compounds  of  chlorine. 

Chlorine  monoxide  is  obtained  by  passing  dry  chlorine  over  red 
oxide  of  mercury  in  the  cold,  and  is  a  pale  yellow,  very  explosive  gas. 
Analysis  shows  that  it  contains  35.46  parts  of  chlorine  to  8  parts  of 
oxygen  by  weight. 

Chlorine  dioxide  is  formed  when  concentrated  sulphuric  acid  is 
added  to  a  small  amount  of  dry  potassium  chlorate  in  a  test-tube, 
and  the  mixture  is  very  cautiously  warmed.  It  is  a  deep  yellow, 
extremely  explosive  gas.  It  contains  35.46  parts  of  chlorine  to  32 
parts  of  oxygen  by  weight. 

A  third  oxide  of  chlorine,  containing  a  higher  proportion  of  oxygen 
than  the  dioxide,  has  also  been  described  (p.  179). 


CHAPTER  IX 

LAWS  OF  CHEMICAL  COMBINATION— 
THE  ATOMIC  THEORY 

Law  of  Constant  Composition  —  In  the  previous  chapters 
we  have  seen  that  the  composition  of  pure  water  is  constant ; 
no  matter  what  its  source  or  how  it  has  been  prepared,  it  always 
contains  88.814  per  cent,  of  oxygen  and  11.186  per  cent,  of  hydrogen 
(approximately  8  parts  by  weight  of  oxygen  to  I  part  of  hydrogen). 
In  the  same  way  we  have  shown  that  hydrogen  chloride  is  of  con- 
stant composition  ;  it  contains  about  35.5  parts  of  chlorine  to  i  part 
of  hydrogen  by  weight.  Further,  it  may  be  shown  that  mercuric 
oxide  always  contains  92.6  per  cent,  by  weight  of  mercury  and 
7.4  per  cent,  by  weight  of  oxygen.  If  oxygen  is  used  in  large 
excess,  nevertheless  the  elements  combine  in  the  above  proportions, 
and  the  excess  of  oxygen  remains  uncombined.  The  composition 
of  an  enormous  number  of  other  pure  substances  has  been  deter- 
mined with  very  great  accuracy,  and  in  all  cases  it  has  been  found 
that  they  are  of  constant  composition.  We  are,  therefore,  justified 
in  assuming  that  this  result  is  a  law  of  nature.  It  is  usually 
called  the  Law  of  Constant  Composition,  and  is  formulated  as 
follows : 

A  definite  chemical  compound  always  contains  the  same  elements 
in  the  same  proportions  by  weight. 

Law  of  Multiple  Proportions— The  determination  of  the  com- 
position of  a  number  of  chemical  compounds  has  shown  that  the  same 
elements  may  unite  in  more  than  one  proportion  to  form  chemical 
compounds.  At  the  end  of  the  last  chapter  it  was  stated  that  there 
are  two  well-defined  oxides  of  chlorine,  which  for  a  fixed  pro- 
portion— 35.5  parts  by  weight — of  chlorine  contain  8  and  32  parts 
of  oxygen  respectively.  Both  are  definite  chemical  compounds  of 
constant  composition,  but  for  a  fixed  proportion  of  chlorine  one 
contains  four  times  as  much  oxygen  as  the  other. 


102     A  TEXT-BOOK  &K  liNORGANIC  CHEMISTRY 


A  fjjitheTr  vll^tratioi)  ^s  iio.-tje  found  in  the  two  oxides  of  lead; 
litharge  or  lead  monoxide,1"  which  'is  yellow,  and  lead  dioxide,  which 
is  dark  brown.  If  weighed  amounts  of  the  two  oxides  are  reduced 
to  metallic  lead  by  heating  in  a  current  of  hydrogen,  it  will  be  found 
that  the  ratio  of  lead  to  oxygen  in  litharge  is  I  :  0.077,  and  in  the 
dioxide  I  :  0.154,  so  that  for  the  same  weight  of  lead  the  ratio  of  the 
amounts  of  oxygen  in  the  two  compounds  is  I  :  2.  Both  are  definite 
chemical  compounds  of  constant  composition,  but  for  the  same 
amount  of  lead  one  contains  twice  as  much  oxygen  as  the  other. 
The  examination  of  a  large  number  of  compounds  shows  that  the 
above  result,  that  for  a  fixed  amount  of  one  element  there  is  a 
simple  relationship  between  the  amounts  of  the  other  element 
present,  is  a  general  rule,  and  we  are  therefore  entitled  to  express 
it  in  the  form  of  a  law.  The  Law  of  Multiple  Proportions,  which 
summarizes  the  facts,  may  be  stated  as  follows  : — 

When  two  elements  unite  in  more  than  one  proportion,  for  a  fixed 
amount  of  one  element  there  is  a  simple  relationship  between  the 
amounts  of  the  other  element  present. 

Later  results  will  show  that  the  ratio  in  question  need  not  be  i  :  2 
or  I  :  4  ;  it  may  be  2  :  3,  3  :  4,  or  even  more  complex.  For  simplicity, 
the  law  has  been  deduced  from  the  composition  of  binary  compounds, 
but  it  applies  to  chemical  compounds  in  general,  whether  simple  or 
complex. 

When  the  composition  of  a  compound  containing  two  elements 
is  expressed  in  percentages,  it  is  not  evident  at  first  sight  that 
the  law  of  multiple  proportions  holds.  Thus  of  the  two  oxides 
of  lead,  the  yellow  compound  contains  92.83  per  cent,  of  lead 
and  7.17  per  cent,  of  oxygen,  and  the  brown  compound  86.6  per 
cent,  of  lead  and  13.4  per  cent,  of  oxygen.  When,  however,  the 
compositions  are  referred  to  a  fixed  amount,  say  I  part,  of  lead, 
the  yellow  compound  is  found  to  contain  0.077  parts,  and  the  brown 
compound  0.154  parts  of  oxygen  to  i  part  of  lead,  the  former 
numbers  being  in  the  ratio  of  i  :  2.  Chlorine  monoxide  contains 
81.59  per  cent,  of  chlorine  and  18.41  per  cent,  of  oxygen  ;  the 
dioxide  contains  52.56  per  cent,  of  chlorine  and  47.44  per  cent,  of 
oxygen.  It  may  readily  be  shown  from  these  figures  that  the  law 
of  multiple  proportions  holds. 

The  Law  of  Combining  Weights — The  study  of  the  com- 
position of  chemical  compounds  has  led  to  the  establishment  of 
a  still  more  comprehensive  law,  of  which  the  two  laws  already 
deduced  are  special  cases.  As  a  preliminary  to  the  deduction  of 


LAWS   OF   CHEMICAL   COMBINATION          103 

the  law,  the  percentage  composition  of  a  number  of  familiar  chemical 
compounds  is  given  in  the  following  table : — 

(1)  Hydrogen  Chloride.        (3)    Water.  (5)  Hydrogen  Sulphide. 

Hydrogen       2.77  Hydrogen  11.18  Hydrogen  5.92 

Chlorine       97.23  Oxygen       88. 81  Sulphur   94.08 

(2)  Magnesium  Chloride.     (4)  Chlorine  Monoxide.     (6)  Sulphur  Dioxide. 

Magnesium  25.54  Chlorine      81.59  Sulphur    50.05 

Chlorine       74.46  Oxygen       18.41  Oxygen    49.95 

The  first  step  in  comparing  the  combining  proportions  of  the  elements 
in  these  six  compounds  is  to  fix  on  one  element  as  a  standard,  to 
which  the  weights  of  the  other  elements  may  be  referred.  As  hydrogen 
appears  to  be  the  element  present  in  smallest  proportion  in  the  com- 
pounds in  question,  we  may  conveniently  take  the  quantity  of  hydrogen 
present  in  a  compound  as  unity,  to  which  the  amounts  of  the  other 
elements  present  are  to  be  referred.  In  our  list  there  are  three  com- 
pounds containing  hydrogen,  and  it  may  easily  be  calculated  that  the 
combining  weight  of  oxygen  in  water  is  approximately  8,  that  of  sulphur 
16,  and  that  of  chlorine  35.5,  when  referred  to  unit  quantity  of  hydrogen. 
•  Taking  now  a  compound  which  contains  two  of  these  elements,  say 
chlorine  monoxide,  we  find  that  35.5  parts  of  chlorine  (the  combining 
weight  of  chlorine  referred  to  hydrogen)  are  combined  with  8  parts  of 
oxygen;  8  is,  however,  the  combining  weight  of  oxygen  as  deter- 
mined by  an  entirely  independent  method — the  analysis  of  water. 
From  this  result  we  might  provisionally  suppose  that  the  proportions 
by  weight  in  which  elements  unite  with  a  definite  amount  of  a  third 
element  are  the  same  as  those  in  which  they  unite  with  each  other.  A 
little  consideration  shows,  however,  that  the  suggested  rule  requires 
amplification.  According  to  the  law  of  multiple  proportions,  two 
elements  may  unite  in  more  than  one  proportion  ;  but,  for  a  fixed 
amount  of  one  element,  there  is  a  simple  ratio  between  the  amounts 
of  the  other  element  present.  Combining  these  results  we  obtain  the 
so-called  Law  of  reciprocal  proportions,  which  may  be  expressed  as 
follows : — 

The  proportions  by  weight  in  which  elements  unite  with  a  fixed 
weight  of  another  element  taken  as  standard  are  the  same,  or  simple' 
multiples  or  submultiples,  of  those  in  which  they  unite  with  each  other. 

The  law  of  reciprocal  proportions,  and  the  other  laws  of  chemical 
combination  already  cited,  are  special  cases  of  a  much  more  compre- 
hensive law,  the  so-called  Law  of  combining  weights,  which  may  be 
formulated  as  follows  : — 

For  each  element  a  fixed  number,  termed  its  combining  weight,  can 


104    A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

be  found,  which  represents  the  ratio,  or  a  simple  integral  multiple  or 
submultiple  of  the  ratio9  in  which  the  element  enters  into  chemical 
combination.  , 

The  proof  of  this  law  follows  at  once  from  a  consideration  of  the 
combining  capacity  of  the  elements  quoted  in  the  table.  The  com- 
bining weight  of  oxygen,  referred  to  hydrogen  as  unit,  is  8,  and  8  is 
also  the  proportion  in  which  oxygen  unites  with  the  combining  weight 
of  chlorine,  35.5.  Further,  the  combining  weight  of  sulphur,  deduced 
from  the  analysis  of  hydrogen  sulphide,  is  16,  and  for  sulphur  dioxide, 
8  parts  of  sulphur  (half  the  combining  weight)  are  combined  with  8 
parts  of  oxygen. 

It  is  evident  from  the  foregoing  that  the  determination  of  the  com- 
bining weight  of  an  element  is  a  comparatively  simple  operation ;  we 
have  only  to  determine  the  proportion  in  which  it  enters  into  com- 
bination with  the  combining  weight  of  another  element.  As  oxygen 
combines  with  more  elements  than  does  hydrogen,  it  has  been  found 
convenient  to  take  the  former  element  as  reference  element,  putting 
its  combining  weight  =  8.  On  this  standard,  the  combining  weight  of 
hydrogen  is  not  strictly  unity,  but  1.008. 

An  interesting  illustration  of  the  determination  of  combining  weights  is  to  be 
found  in  the  action  of  different  metals  on  an  acid  such  as  hydrochloric  acid. 
Known  weights  of  zinc,  magnesium,  iron,  and  aluminium  are  treated  with  excess 
of  dilute  hydrochloric  acid  (cf.  p.  126),  and  the  volume  of  hydrogen  measured  in 
each  case.  The  relative  amounts  of  the  different  metals  which  liberate  the  same 
volume  or  weight  of  hydrogen  (say  i  part  by  weight  of  hydrogen)  are  in  the  ratio 
Zn=32-5,  Mg=i2.i8,  Fe=28,  Al=7,  and  this  is  necessarily  the  ratio  of  the  com- 
bining weights,  since  it  represents  the  relative  weights  of  the  different  metals 
which  have  entered  into  chemical  combination  with  equal  weights  of  chlorine  *  (as 
already  pointed  out,  there  is  a  definite  and  invariable  ratio  between  the  amounts 
of  chlorine  and  hydrogen  in  hydrogen  chloride).  These  relative  amounts  of  the 
different  metals  are  also  chemically  equivalent,  since  they  set  free  equal  volumes  of 
hydrogen,  and  therefore  combining  weights  are  often  termed  chemical  equivalents. 

We  may  therefore,  from  an  experimental  standpoint,  define  the 
combining  weight  or  chemical  equivalent  of  an  element  as  that 
amount  of  it  which  can  combine  with  (or  take  the  place  of)  8  parts  by 
weight  of  oxygen  or  1.008  parts  by  weight  of  hydrogen.  We  shall 
learn  more  fully  in  the  next  chapter  that  an  element  may  have  more 
than  one  chemical  equivalent. 

Gay-Lussac's  Law  of  Volumes  —Having  discussed  the  laws 
of  combination  by  weight,  we  now  proceed  to  consider  the  laws  of 

1  The  combining  weight  of  magnesium  as  given  here  may  be  compared  with 
that  calculated  from  the  composition  of  magnesium  chloride  (see  table,  p.  103). 


THE   ATOMIC   THEORY  105 

combination  by  volume  ;  and  in  this  connection  we  confine  ourselves 
to  the  volumes  of  gases  and  vapours,  as  offering  the  simplest  relation- 
ships. It  has  been  shown  that  hydrogen  and  oxygen  combine  in  the 
ratio  of  two  volumes  of  the  former  to  one  volume  of  the  latter,  giving 
two  volumes  of  steam,  when  all  the  gases  are  measured  under  the 
same  conditions.  Further,  one  volume  of  hydrogen  combines  with 
one  volume  of  chlorine  to  form  two  volumes  of  hydrogen  chloride. 
These  remarkably  simple  relationships,  and  others  of  a  similar 
kind,  were  discovered  by  the  French  chemist  Gay-Lussac,  and  led 
him  to  the  enunciation  of  the  Law  of  Gaseous  Volumes,  which  may 
be  stated  as  follows  : — 

Gases  combine  in  simple  ratios  by  volume,  and  the  volume  of  the 
gaseous  product  bears  a  simple  ratio  to  the  volumes  of  the  reacting 
gases. 

Subsequent  investigation  has  shown  that  the  Law  of  Volumes,  like 
the  other  gas  laws,  is  very  nearly,  but  not  exactly  true. 

The  Atomic  Theory— The  four  laws  of  chemical  combination 
just  considered  are  purely  experimental,  and  independent  of  any  view 
we  may  take  as  to  the  constitution  of  matter.  In  connexion  with 
the  properties  of  gases,  however,  we  have  already  discussed  the  theory 
that  matter  is  not  continuous,  but  is  made  up  of  extremely  small, 
practically  incompressible  particles,  which  are  far  apart  in  gases  but 
much  closer  together  in  solids  and  liquids,  and  it  is  natural  to  inquire  if 
this  theory  throws  any  light  on  the  Jaws  of  chemical  combination. 

The  view  of  the  constitution  of  matter  just  indicated  originated 
with  the  Greek  philosophers,  but  was  first  developed  to  a  consistent 
theory  by  Dalton  (1808)  ;  it  is  termed  the  atomic  theory.  According 
to  this  theory,  matter  is  made  up  of  small  particles  called  atoms, 
which  cannot  be  further  divided  by  any  means  at  our  disposal.  The 
atoms  of  any  one  element  are  identical  in  all  respects,  but  differ,  at 
least  in  weight,  from  those  of  other  elements.  By  the  union  of  the 
atoms  of  different  elements  in  simple  numerical  proportions,  chemical 
compounds  are  formed. 

As  was  first  pointed  out  by  Dalton,  the  laws  of  chemical  combina- 
tion by  weight  find  a  ready  explanation  on  the  atomic  theory.  For 
simplicity  we  will  consider  only  binary  compounds,  that  is,  compounds 
made  up  of  two  elements  ;  but  the  reasoning  is  the  same  for  more 
complicated  compounds.  As  the  ultimate  particles  of  a  binary  com- 
pound are  made  up  of  two  kinds  of  atoms,  and  a  fixed  number  of  each 
kind,  the  compound  must  be  of  constant  composition,  since  it  is  made 
up  of  an  enormous  number  of  such  ultimate  particles.  If,  for  example, 


io6    A   TEXT-BOOK  OF   INORGANIC   CHEMISTRY 

we  assume  that  the  ultimate  particles  of  hydrogen  chloride  are  made 
up  of  one  atom  of  hydrogen,  weight  1.008,  and  one  atom  of  chlorine, 
weight  35.46,  the  compound  must  be  of  invariable  composition,  con- 
taining hydrogen  and  chlorine  in  the  ratio  1.008  :  35.46. 

Similarly,  the  law  of  multiple  proportions  follows  at  once  from  the 
atomic  theory.  If  an  ultimate  particle  of  a  compound  contains  an 
atom  of  one  element,  of  weight  .r,  and  an  atom  of  another  element,  of 
weight  y,  the  compound  will  contain  the  two  elements  in  the  ratio 
x  \y.  If  the  same  two  elements  unite  to  form  a  second  compound, 
the  ultimate  particles  of  which  contain  one  atom  of  the  first  element 
and  two  of  the  second,  the  ratio  of  the  two  elements  in  the  second 
compound  will  be  x  :  2.y.  From  this  it  follows  that  for  a  fixed  amount 
x  of  one  element,  the  other  element  is  present  in  the  two  compounds 
in  the  exact  ratio  1:2,  in  accordance  with  the  law  of  multiple 
proportions. 

Finally,  the  law  of  combining  weights  is  also  seen  to  be  a  logical 
consequence  of  the  theory,  the  experimentally  found  combining 
weights  bearing  a  simple  relation  to  the  relative  weights  of  the  atoms. 
The  fixing  of  the  relative  weights  of  the  atoms  is  clearly  a  matter  of 
the  utmost  importance,  and  we  shall  see  in  the  following  paragraphs 
how  the  problem  was  satisfactorily  solved. 

Avogadro's  Hypothesis— We  have  seen  that,  according  to 
Gay-Lussac's  law,  gases  combine  in  simple  ratios  by  volume.  Further, 
according  to  the  atomic  theory,  chemical  combination  takes  place 
between  one  or  two  (or  some  small  number  of)  atoms  of  one  kind  and 
one  or  two  atoms  of  another  kind.  By  combining  these  two  state- 
ments it  is  evident  that  there  is  a  simple  relationship  between  the 
number  of  particles  in  equal  volumes  of  different  gases  under  the  same 
conditions.  Strong  support  is  lent  to  this  deduction  by  the  facts 
already  stated  with  reference  to  the  similar  behaviour  of  all  gases  on 
altering  the  temperature  or  pressure. 

The  most  obvious  suggestion  in  this  connexion  is  that  equal  volumes 
of  all  gases  contain  the  same  number  of  atoms,  and  Gay-Lussac  him- 
self was  at  first  inclined  to  adopt  this  view.  It  was  soon  found, 
however,  to  be  untenable  ; *  and  the  view  hsld  at  the  present  day  was 

1  In  order  to  understand  the  difficulties  experienced  by  Dalton  and  by  Gay- 
Lussac  in  reconciling  the  law  of  volumes  with  the  atomic  theory,  it  must  be 
remembered  that  Dalton  applied  the  term  atom  to  chemical  compounds  as  well 
as  to  elements,  meaning  thereby  the  smallest  particles  capable  of  independent 
existence.  Thus  the  atom  of  hydrogen  chloride  was  looked  upon  as  being  formed 
by  the  combination  of  an  atom  of  hydrogen  and  an  atom  of  chlorine.  If  now 
equal  volumes  contain  the  same  number  of  atoms,  one  volume  of  hydrogen  should 


THE   ATOMIC   THEORY  107 

shortly  afterwards  put  forward  by  Avogadro  (1811).  Avogadro  drew 
a  distinction  between  atoms,  the  smallest  particles  of  matter  which 
can  take  part  in  chemical  changes,  and  molecules,  the  smallest  par- 
ticles of  matter  capable  of  independent  existence,  and  formulated  his 
hypothesis  as  follows  : — 

Equal  volumes  of  all  gases,  under  the  same  conditions  of  tempera- 
ture and  pressure,  contain  the  same  number  of  molecules. 

There  should  be  no  difficulty  in  drawing  a  clear  distinction  between 
molecules  and  atoms.  Atoms,  as  the  name  implies,  are  not  divisible 
by  any  means  at  our  disposal ;  molecules,  on  the  other  hand,  are 
almost  invariably  complex,  being  made  up  of  two  or  more  atoms. 
The  molecules  of  a  chemical  compound,  the  smallest  particles  capable 
of  independent  existence,  are  made  up  of  atoms  of  different  kinds. 
On  the  other  hand,  the  molecules  of  which  a  quantity  of  an  elementary 
substance  is  made  up,  are  usually  formed  by  the  union  of  two  or 
more  atoms  of  an  element ;  in  a  few  cases  by  a  single  atom  of  the 
element.  We  shall  see  later  that  the  atoms  of  elements  already 
considered — hydrogen,  oxygen,  chlorine — have  not  been  obtained 
independently  ;  the  molecule  of  hydrogen,  for  example,  is  formed 
by  the  union  of  two  atoms  of  hydrogen.  On  the  other  hand,  the 
molecule  of  mercury  in  the  gaseous  state  is  made  up  of  a  single  atom. 

It  will  be  shown  later  that  Avogadro's  hypothesis  leads  to  the  con- 
clusion that  the  molecule  of  hydrogen  chloride  is  made  up  of  one  atom 
of  hydrogen  and  one  atom  of  chlorine.  Assuming  this  result  for  the 

combine  with  one  volume  of  chlorine  to  form  one  volume  of  hydrogen  chloride, 
which  is  contrary  to  the  facts.  The  matter  becomes  clear,  however,  when  with 
Avogadro  we  make  the  assumption  that  the  smallest  independent  particles  of 
elementary  gases  are  not  indivisible,  as  Dalton  suggested,  but  in  the  great 
majority  of  cases  are  composed  of  two  or  more  atoms.  Avogadro's  reasoning  was 
somewhat  as  follows  (cf.  p.  108).  If  the  hydrogen  chloride  formed  by  combina- 
tion of  equal  volumes  of  hydrogen  and  chlorine  without  contraction  contains  a 
number  of  molecules  equal  to  the  sum  of  the  number  of  molecules  of  hydrogen 
and  chlorine  before  combination,  then  the  compound  cannot  be  formed  by  an 
association  of  previously  separate  particles,  which  would  lead  to  a  diminution  in 
the  number  of  particles.  On  the  contrary,  the  molecules  of  hydrogen  and 
chlorine  must  be  complex,  and  an  exchange  of  ultimate  particles  (atoms)  be- 
tween the  molecules  must  occur  ;  that  is,  the  hydrogen  and  chlorine  molecules 
must  each  split  into  (two)  atoms,  the  atoms  of  different  kinds  then  uniting  in  pairs 
to  form  new  complex  molecules.  Expressing  the  result  graphically  we  have 


H 


H     .+..    Cl 


the  number  of  particles  thus  remaining  unchanged. 


io8     A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 


moment,  the  accompanying  diagram  gives  a  graphic  representation 
of  Avogadro's  hypothesis  applied  to  hydrogen  chloride,  hydrogen, 
and  mercury.  The  symbol  @  represents  an  atom  of  hydrogen, 
@  represents  an  atom  of  chlorine,  and  (Hg)  represents  an  atom  of 
mercury.  The  spaces  a,  bt  and  c  represent  equal  volumes,  and 
the  diagram  illustrates  the  following  important  points  :  (i)  Equal 
volumes  of  all  gases  and  vapours  contain  the  same  number  of  mole- 
cules under  the  same  conditions,  no  matter  whether  the  molecules 
are  those  of  elements  or  of  chemical  compounds.  (2)  The  volumes  of 
the  molecules  themselves  (the  "  particles"  of  the  kinetic  theory,  p.  48) 


®o 


©     © 
©     © 


© 


© 


©     ©     © 
®     ©     © 


®@  ®@ 


©© 
®@ 


®@ 


@©  ®® 


®@ 


are  small  compared  with  the  total  volume  occupied.  (3)  The  mole- 
cules of  chemical  compounds  are  made  up  of  atoms  of  different  kinds. 
(4)  The  molecules  of  elements  are  made  up  of  atoms  of  the  same  kind, 
and  the  number  of  atoms  in  the  molecule  varies  from  one  upwards. 

These  four  statements  contain  the  essential  features  of  the  mole- 
cular theory  of  matter.  The  second  statement  at  once  enables  us  to 
understand  how  the  same  number  of  molecules  of  different  substances, 
although  they  may  be  very  different  in  size,  may  still  occupy  equal 
volumes  in  the  gaseous  form  under  equivalent  conditions. 

Complexity  of  the  Molecules  of  Hydrogen  and  of 
Oxygen — We  have  seen  that  one  volume  of  hydrogen  and  one 
volume  of  chlorine  unite  to  form  two  volumes  of  hydrogen  chloride. 
Suppose  in  these  two  volumes  there  are  2000  molecules  of  hydrogen 


THE   ATOMIC   THEORY  109 

chloride.  Each  of  these  molecules  contains  some  hydrogen  and 
chlorine,  and  must  contain  at  least  one  atom  of  each.  Therefore, 
2000  atoms  of  hydrogen  and  2000  atoms  of  chlorine  must  have  been 
concerned  in  the  formation  of  the  2000  molecules  of  hydrogen  chloride. 
But,  according  to  Avogadro's  hypothesis,  if  two  volumes  of  the  hydro- 
gen chloride  contain  2000  molecules,  the  one  volume  of  hydrogen 
must  contain  1000  molecules.  It  follows,  therefore,  that  the  2000 
atoms  of  hydrogen  have  been  obtained  from  1000  molecules  of  the 
gas,  so  that  each  molecule  of  the  gas  must  contain  (at  least)  two 
atoms.  By  exactly  similar  reasoning  it  may  be  shown  that  each 
molecule  of  chlorine  is  made  up  of  two  atoms  of  that  element. 

It  is  evident  that  the  above  reasoning  only  fixes  a  lower  limit  to  the 
complexity  of  the  hydrogen  molecule ;  if  the  molecule  of  hydrogen 
chloride  contains  more  than  one  atom  of  hydrogen,  the  hydrogen 
molecule  must  be  still  more  complex.  At  a  later  stage,  however 
(p.  1 16),  evidence  will  be  given  that  hydrogen  chloride  contains  only 
one  atom  of  each  element,  and  therefore  the  molecules  of  hydrogen 
and  of  chlorine  contain  only  two  atoms  of  the  respective  elements. 

From  the  experimental  fact  that  two  volumes  of  hydrogen  and  one 
volume  of  oxygen  give  two  volumes  of  steam,  it  may  be  shown  in 
exactly  the  same  way  that  the  oxygen  molecule  is  formed  by  the 
union  of  two  atoms. 

Avogadro's  Hypothesis  and  Molecular  Weights— On 
the  basis  of  Avogadro's  hypothesis,  the  determination  of  the  relative 
weights  of  the  molecules  of  gases  is  very  simple.  As  equal  volumes 
contain  the  same  number  of  molecules,  it  is  evident  that  the  relative 
weights  of  equal  volumes  of  gases  give  us  the  relative  weights  of  the 
molecules  contained  in  them  ;  in  other  words,  the  ratio  of  the  densities 
of  two  gases  is  the  ratio  of  the  weights  of  their  molecules.  It  only 
remains  to  choose  some  substance  as  standard,  to  which  the  weights 
of  the  molecules  of  other  substances  are  to  be  referred.  As  hydrogen 
is  the  lightest  gas,  and  has  already  been  used  as  standard  in  density 
determinations,  it  is  natural  to  use  it  also  as  standard  for  molecular 
weights.  On  this  basis  the  weight  of  the  molecule  of  hydrogen 
might  be  taken  as  unity.  It  has  been  pointed  out,  however,  that  the 
hydrogen  molecule  contains  two  atoms,  and  there  are  certain  advan- 
tages in  taking  the  weight  of  the  atom  of  hydrogen  as  unity,  to  which 
all  molecular  weights  are  referred.  On  this  basis,  the  molecular 
weight  of  hydrogen  is  2. 

In  previous  chapters  it  has  been  pointed  out  that  the  density  of 
oxygen  in  round  numbers  is  16,  referred  to  hydrogen  as  standard,  and 


no     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

this,  according  to  Avogadro's  hypothesis,  indicates  that  the  oxygen 
molecule  is  sixteen  times  the  weight  of  the  hydrogen  molecule,  or 
thirty-two  times  the  weight  of  the  hydrogen  atom.  The  molecular 
weight  of  oxygen  is  therefore  found  by  doubling  its  vapour  density. 
This  method  is  clearly  of  general  application,  and  may  be  stated  as 
follows : — 

The  molecular  'weight  of  a  gas  is  double  its  vapour  density  referred 
to  hydrogen  as  unit. 

In  order  that  the  basis  of  this  important  rule  may  be  clearly 
realized,  we  may  think  of  a  certain  distance  which,  when  expressed 
in  feet,  measures  15,  and  in  yards  measures  5.  It  is  clear  that,  since 
in  the  first  case  the  distance  is  referred  to  a  unit  one-third  of  that 
used  in  the  second  case,  the  number  expressing  the  distance  in  the 
first  case  must  be  three  times  that  in  the  second  case.  In  exactly  the 
same  way,  molecular  weights,  being  referred  to  a  standard  half  that 
to  which  the  densities  are  referred,  must  be  represented  by  numbers 
which  are  double  those  representing  the  densities.  A  slightly  different 
way  of  stating  Avogadro's  hypothesis  is  that  the  same  volume  is 
occupied x  by  the  molecules  of  all  gases  under  the  same  conditions ; 
in  other  words,  the  molecular  volumes  of  all  gases  are  equal.  It 
is  found  convenient  to  represent  each  molecule  as  occupying  unit 
volume,  a  convention  which  will  be  largely  used  later  on.  No 
assumption  is  made  as  to  the  actual  volume  occupied  by  a  molecule, 
the  definition  being  only  used  relatively. 

It  can  readily  be  calculated  from  the  density  given  on  p.  36,  that 
the  molecular  weight  of  hydrogen  in  grams  (that  is,  2.016  grams  of 
hydrogen)  occupies  about  22.40  litres  at  o°  and  760  mm.  pressure.  As, 
however,  the  molecules  of  all  gases  occupy  the  same  volume  under 
the  same  conditions,  the  molecular  weight  of  any  other  gas  in  grams 
must  also  occupy  22.40  litres  under  normal  conditions.  This  is  the 
definition  of  molecular  weight  which  should  be  remembered  by  the 
student^  and  it  will  be  repeated  for  the  sake  of  definiteness. 

The  molecular  weight  of  a  gas  is  that  weight  of  it,  expressed  in 
gramS)  which  occupies  22.4  litres  at  o°  and  760  mm.  pressure. 

The  validity  of  the  above  statement  may  perhaps  be  shown  still  more 

1  The  term  occupied  is  here  used  in  the  sense  of  inhabited,  and  does  not  mean 
that  the  space  infilled  with  the  particles.  The  true  meaning  is  rendered  evident 
by  the  diagram  on  p.  107.  In  equal  volumes  of  different  gases  there  are  the 
same  number  of  particles,  therefore  single  molecules  of  different  gases  inhabit 
the  same  average  space  under  corresponding  conditions.  For  purposes  of  com- 
parison we  take  this  average  space  as  unit  volume. 


THE  ATOMIC   THEORY  in 

clearly  by  taking  it  in  the  converse  way.  In  22.4  litres  of  different 
gases  there  are  the  same  number  of  molecules,  according  to  Avogadro's 
hypothesis,  and  the  relative  weights  of  these  equal  volumes  are  clearly 
in  the  ratio  of  the  molecular  weights.  That  particular  volume,  how- 
ever, contains  the  molecular  weight  of  hydrogen  in  grams,  and 
therefore  the  molecular  weights  of  the  other  gases  in  grams. 

The  enormous  importance  of  the  statements  just  made  is  obvious. 
If  a  substance  can  be  vaporized  without  decomposition,  its  molecular 
weight  can  be  determined  without  difficulty.  We  shall  see  later 
that  other  methods  of  determining  molecular  weights  are  used  for 
substances  which  cannot  be  converted  into  vapour  without  decom- 
position (cf.  Chapter  XV.). 

Symbols  and  Formulae  —  Before  proceeding  to  discuss  the 
methods  used  in  determining  atomic  weights,  it  is  desirable  to  con- 
sider briefly  the  methods  used  in  chemistry  for  representing  the  com- 
position of  elements  and  compounds.  Instead  of  writing  out  in  full  the 
name  of  an  element  each  time  it  is  mentioned,  it  is  conveniently  and 
shortly  indicated  by  using  the  first  letter,  or,  in  certain  cases,  two 
letters  of  its  Latin  name.  Thus,  O  represents  oxygen,  N  stands  for 
nitrogen,  C  for  carbon,  S  for  sulphur,  and  so  on.  The  symbol  for 
chlorine  is  Cl ;  for  tin  Sn  (from  stannum) ;  for  lead  Pb  (plumbum)  ;  for 
iron  Fe  (contracted  from  ferrum) ;  and  antimony  is  represented  by  Sb 
(stibium).  When  the  names  of  more  than  one  element  begin  with  the 
same  letter,  the  initial  letter  is  generally  used  for  the  better  known 
element,  and  as  symbols  for  the  other  elements  two  letters  are  used. 
A  full  list  of  the  symbols  for  the  elements  is  given  on  the  back  page 
of  the  cover,  and  the  student  should  gradually  familiarize  himself 
with  them. 

Chemical  compounds  are  formed  by  the  combination  of  two  or 
more  elements,  and  it  is  natural  to  indicate  the  composition  of  a  com- 
pound by  writing  side  by  side  the  symbols  of  the  elements  composing 
it.  It  will,  however,  obviously  be  an  enormous  advantage  if  the 
symbols  can  be  used  to  indicate,  not  only  the  qualitative  but  also  the 
quantitative  composition  of  a  chemical  compound.  The  first  step 
towards  this  end  is  to  use  the  symbol  of  an  element,  not  merely  to 
indicate  the  presence  of  the  element,  but  to  represent  the  atomic 
weight  of  the  element.  Thus  the  symbol  H,  when  used  to  represent 
a  component  of  a  chemical  compound,  stands  for  i  part  by  weight  of 
hydrogen,  and  Cl  stands  for  35.46  parts  of  chlorine.  As  already 
indicated,  the  molecule  of  hydrogen  chloride  is  formed  by  the  union 
of  one  atom  of  hydrogen  and  one  atom  of  chlorine  This  is  indicated 


ii2     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

by  HC1,  which  shows  not  only  that  hydrogen  chloride  contains 
hydrogen  and  chlorine,  but  that  it  contains  i  part  by  weight  of 
hydrogen  to  35.46  parts  by  weight  of  chlorine.  HC1  is  termed  the 
formula  for  hydrogen  chloride.  In  the  same  way,  the  formula  for 
common  salt,  NaCl,  indicates  that  it  contains  sodium  and  chlorine, 
and  that  these  elements  are  present  in  the  ratio  of  23  parts  by  weight 
of  sodium  to  35.46  parts  by  weight  of  chlorine. 

It  sometimes  happens  that  the  molecule  of  a  chemical  compound 
contains  more  than  one  atom  of  a  particular  element.  The  number 
of  atoms  present  is  indicated  by  the  appropriate  figure  written  at  the 
lower  right-hand  side  of  the  symbol.  Thus  H2O,  which  is  the  formula 
for  water,  indicates  that  the  molecule  contains  two  atoms  of  hydrogen 
to  one  atom  of  oxygen,  and  as  the  atomic  weights  of  these  elements 
are  in  the  approximate  ratio  of  i  :  16,  the  formula  indicates  that  water 
is  made  up  of  2  parts  by  weight  of  hydrogen  to  16  parts  by  weight 
of  oxygen,  as  has  already  been  proved  experimentally.  The  remark- 
able simplicity  and  convenience  of  this  system  of  formulation  will  be 
clear  from  the  example  just  given. 

From  the  formula  of  ordinary  chalk,  CaCO3,  it  may  easily  be  calcu- 
lated that  it  contains  40  per  cent,  of  calcium,  12  per  cent,  of  carbon, 
and  48  per  cent,  of  oxygen.  The  composition  of  potassium  chlorate, 
KC1O3,  and  of  sulphuric  acid,  H2SO4,  may  also  be  calculated  from  the 
formulae.1 

In  some  cases  a  number  of  chemical  compounds  contain,  the  same 
group  of  elements.  We  have  just  seen  that  the  formula  for  potassium 
chlorate  is  KC1O3,  and  it  will  be  shown  later  that  all  chlorates  contain 
the  group  C1O3  ;  that  is,  they  contain  chlorine  and  oxygen  in  the 
proportion  of  one  atom  of  the  former  to  three  atoms  of  the  latter. 
Sometimes,  however,  a  molecule  contains  more  than  one  such  group 
of  elements  ;  for  example,  barium  chlorate  contains  two  C1O3  groups 
associated  with  one  atom  of  barium.  This  might,  of  course,  be 
indicated  by  writing  its  formula  BaCl2O6,  but  this  method  has  the 
disadvantage  that  we  could  not  tell  at  a  glance  that  we  are  dealing 
with  a  chlorate.  It  is  therefore  preferable  to  write  the  formula  of 
barium  chlorate  either  as  Ba(ClO3)2  or  as  Ba2ClO3.  An  integer  on 
the  line,  as  in  the  last  example,  is  to  be  understood  as  multiplying 
all  that  follows,  up  to  the  next  comma. 

Still  another  type  of  formula  may  be  mentioned  here  for  the  sake 
of  completeness.    We  have  already  seen  that  some  salts  form  chemical 
1  Approximate   atomic   weights:    H  =  i;    €  =  12;   O=i6;    8=32;    €1  =  35.5. 
K=39. 


THE   ATOMIC   THEORY  113 

compounds  with  water,  generally  called  hydrates.  Thus  when  sodium 
chloride  separates  from  solution  at  low  temperatures,  one  molecule 
of  it  is  found  to  be  combined  with  two  molecules  of  water.  When 
the  temperature  is  raised,  the  two  molecules  of  water  are  driven  off 
and  sodium  chloride  is  left.  The  compound  in  question  is  repre- 
sented by  the  formula  NaCl52H2O.  When  two  parts  of  a  formula  are 
thus  separated  by  a  comma,  it  usually  indicates  that  the  compound 
is  comparatively  easily  broken  up  at  the  point  in  question.  The 
water  associated  with  such  compounds  is  sometimes  called  water  of 
crystallization. 

The  meaning  of  the  formula  C12,8H2O,  ascribed  to  chlorine  hydrate, 
will  now  be  readily  understood. 

On  the  basis  of  Avogadro's  hypothesis,  we  have  decided  to  use  the 
convention  that  the  molecule  of  any  substance  in  the  state  of  vapour 
occupies  unit  volume.  The  formulae  of  the  volatile  substances 
quoted  in  the  present  section,  since  they  represent  the  formulae  of 
molecules,  are  termed  molecular  formula ',  and  stand  for  the  amounts  of 
the  substances  which  occupy  unit  volume  in  the  form  of  vapour.  A 
formula  such  as  HC1  is  therefore  remarkably  expressive,  inasmuch  as — 

(1)  It  shows  what  elements  are  present  in  the  compound. 

(2)  It  shows  the  relative  proportions  by  weight  of  the  components. 

(3)  It  represents  unit  volume  of  the  substance  in  the  state  of  vapour. 
We  have    now  met  with   two   sets   of  numbers,   the   combining 

weights  and  the  atomic  weights,  both  of  which  represent  the  propor- 
tions in  which  elements  enter  into  chemical  combination.  The  exact 
relationship  between  these  two  sets  of  numbers,  and  the  methods 
employed  in  fixing  atomic  weights,  form  the  subject  matter  of  the 
next  chapter. 

Fact— Generalization  or  Natural  Law— Hypothesis- 
Theory — It  will  be  convenient  to  conclude  this  chapter  with  a  brief 
account  of  the  theoretical  principles  underlying  scientific  investigation, 
as  illustrated  in  the  foregoing  pages. 

Chemistry,  like  all  the  experimental  sciences,  is  based  on  facts, 
established  by  experiment.  A  large  number  of  such  facts  have 
already  been  mentioned  ;  for  example,  that  certain  chemical  com- 
pounds, which  have  been  investigated  with  the  greatest  care,  always 
contain  the  same  elements  in  the  same  proportions.  A  mere  collec- 
tion of  facts,  however,  does  not  constitute  a  science.  When  a  certain 
number  of  facts  have  been  established,  the  chemist  proceeds  to 
reason  from  analogy  as  to  the  behaviour  of  systems  under  conditions 
which  have  not  yet  been  investigated.  For  example,  Proust  showed 
8 


ii4     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

that  there  are  two  well-defined  oxides  of  tin  and  that  the  composition 
of  each  is  invariable.  From  the  results  of  these  and  a  few  other 
investigations  he  concluded  from  analogy  that  the  composition  of  all 
pure  chemical  compounds  is  invariable,  although,  of  course,  very  few 
of  them  had  then  been  investigated  from  that  point  of  view.  To 
proceed  in  this  way  is  to  generalize,  and  the  short  statement  of  the 
conclusion  arrived  at  is  termed  a  generalization  or  law.  It  will  be 
evident  that  a  law  is  not  in  the  nature  of  an  absolute  certainty ;  it 
comprises  the  facts  experimentally  established,  but  also  enables  us  to 
foretell  a  great  many  things  which  have  not  been,  but  which  if  necessary 
could  be  investigated  experimentally.  The  greater  the  number  of 
cases  in  which  a  law  has  been  found  to  hold,  the  greater  is  the 
confidence  in  its  validity,  until  finally  a  law  may  attain  practically 
the  same  standing  as  a  statement  of  facts.  We  may  confidently  antici- 
pate that,  however  greatly  our  views  regarding  natural  phenomena 
may  change,  such  generalizations  as  the  law  of  constant  proportions 
will  remain  eternally  true. 

Natural  laws  can  be  discovered  in  two  ways:  (i)  by  correlating  a 
number  of  experimental  facts,  as  just  indicated:  this  is  termed 
induction;  (2)  by  a  speculative  method,  on  the  basis  of  certain 
hypotheses  as  to  the  nature  of  the  phenomena  in  question,  the  con- 
sequences of  which  are  then  tested  by  experiment :  this  is  termed 
deduction.  The  meaning  to  be  attached  to  the  term  hypothesis  is 
best  illustrated  by  an  example.  In  the  present  chapter,  we  have 
seen  that  the  laws  of  chemical  combination  are  accounted  for  satis- 
factorily on  the  view  that  matter  is  made  up  of  extremely  small,  dis- 
crete particles,  the  atoms.  Such  a  mechanical  representation,  which 
is  more  or  less  inaccessible  to  experimental  proof,  is  termed  a  hypo- 
thesis. A  hypothesis  may,  therefore,  be  defined  as  a  mental  picture 
of  an  unknown,  or  largely  unknown,  state  of  affairs,  in  terms  of 
something  which  is  better  known.  Thus  the  state  of  affairs  in  gases, 
which  is  and  will  remain  unknown  to  us,  is  represented,  according 
to  the  kinetic  theory,  in  terms  of  an  enormous  number  of  rapidly 
moving,  perfectly  elastic  particles,  and  on  this  basis  it  is  possible  to 
deduce  certain  of  the  laws  which  are  actually  followed  by  gases  (p.  49). 

There  does  not  appear  to  be  any  fundamental  distinction  in  the 
use  of  the  terms  "hypothesis"  and  "theory."  A  theory  may  be 
defined  as  a  hypothesis,  many  of  the  deductions  from  which  have 
been  confirmed  by  experiment  and  which  admits  of  the  convenient 
representation  of  a  large  number  of  experimental  facts. 


CHAPTER  X 

DETERMINATION  OF  ATOMIC  WEIGHTS— COMBIN- 
ING WEIGHTS  AND  CHEMICAL  EQUIVALENTS- 
FORMULAE  AND  EQUATIONS— VALENCY 

IN  Chapter  IV.  we  have  been  led  to  assign  definite  combining 
weights  to  each  element,  and  a  few  of  these  combining  weights 
are  given  in  the  accompanying  table.  The  combining  weight  of  an 
element  has  already  been  defined  as  the  smallest  quantity  of  it  which 
combines  with,  or  takes  the  place  of,  I  part  by  weight  of  hydro- 
gen, or  8  parts  by  weight  of  oxygen.  On  the  other  hand,  we  have 
made  use  in  the  last  chapter  of  numbers  called  atomic  weights  ; 
these  are  supposed  to  represent  the  relative  weights  of  the  atoms  of 
other  elements,  referred  to  hydrogen  as  unity,  more  accurately,  to 
oxygen  =  16,  as  in  the  case  of  the  combining  weights. 

A  comparison  of  the  experimentally  found  combining  weights,  and 
of  the  atomic  weights,  for  a  number  of  elements  is  shown  in  the 
accompanying  table : — 

Hydrogen.     Oxygen.     Sulphur.     Chlorine.     Magnesium. 
Combining  Weights      .     1.008  8.0  16.0  35-46  12.16 

Atomic  Weights    .         .     1.008  16.0  32.0  35-46  24.32 

It  is  evident  from  these  data  that  the  atomic  weights  are  either  the 
same  as,  or  are  simple  integral  multiples  of,  the  combining  weights. 
There  must  clearly  be  some  good  reason  why  atomic  weights  are 
preferred  to  combining  weights  in  representing  chemical  formulas, 
and  these  will  be  fully  set  forth  in  the  following  sections  (p.  120). 

Determination  of  Atomic  Weights  by  Volumetric 
Methods — The  molecular  weight  of  a  volatile  substance — the 
weight  of  the  molecule  referred  to  the  atom  of  hydrogen  as  unit — 
can  be  determined  on  the  basis  of  Avogadro's  hypothesis  ;  it  is  the 
amount  of  the  substance  in  grams  which,  at  o°  and  760  mm. 
pressure,  occupies  22.4  litres.  The  molecule  is  made  up  of  atoms, 
and  the  molecular  weight  is  the  sum  of  the  weights  of  the  atoms. 
The  amount  contributed  to  the  molecular  weight  by  each  of  the 

"3 


1 6     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


atoms  can  of  course  readily  be  determined  by  analysis.  Experiment 
shows  that  the  molecular  weight  of  hydrogen  chloride  is  36.46 
approximately,  and  an  analysis  of  the  compound  shows  that  about 
i  part  is  hydr6gen  and  the  remainder,  35.46,  is  chlorine.  If  we 
know  further  how  many  atoms  of  hydrogen  and  of  chlorine  are 
present  in  the  molecule,  we  at  once  obtain  the  atomic  weight. 
If,  for  example,  only  one  atom  of  chlorine  is  present,  the  atomic 
weight  of  this  element  must  be  35.46,  if  two  atoms  are  present  it 
must  be  17.73.  The  problem  is  solved  if  we  know  how  many  atoms  of 
an  element  are  present,  but  this  is  just  where  the  difficulty  comes  in. 
It  may  safely  be  assumed,  however,  that  if  we  investigate  a  sufficient 
number  of  volatile  compounds,  we  shall  meet  with  some  which  con- 
tain not  more  than  one  atom  in  the  molecule.  The  relative  amount 
of  the  element  in  those  molecules  in  which  it  occurs  in  the  smallest 
proportion  is  therefore  the  atomic  weight  required. 

The  atomic  weight  of  an  element  is  therefore  the  smallest  amount 
of  it  which  occurs  in  a  molecule  of  one  of  its  compounds,  the  unit  of 
weight  being  ^  of  the  atomic  weight  of  oxygen  (cf.  p.  104). 

As  an  illustration  of  the  method,  we  will  use  it  to  find  the  atomic 
weights  of  hydrogen,  oxygen,  and  chlorine,  elements  which  have 
been  considered  in  the  previous  chapters,  and  also  of  carbon  and 
mercury. 


Weights  of  Constituents  in  Molecular  Weights. 

Compound. 

Molecular 
Weights. 

Hydrogen. 

Oxygen. 

Chlorine. 

Carbon. 

Mercury. 

Hydrogen  chloride 

36.S 

I 

35-5 

Chlorine  monoxide 

8? 

... 

io 

71 

... 

0> 

Chlorine  dioxide 

67-5 

... 

32 

35-5 

... 

t> 

Water      .     .     . 

18 

2 

16 

... 

M 

Methane  . 

16 

4 

M 

12 

Ethylene  .     .     . 

28 

4 

M 

24 

Acetylene     .     . 

26 

2 

... 

.. 

24 

.. 

Carbon  monoxide 

28 

... 

16 

M 

12 

... 

Carbon  dioxide 

44 

... 

32 

M 

12 

... 

Mercurous  chloride 

23S-5 

... 

35-5 

... 

2OO 

Mercuric  chloride  . 

271 

... 

... 

7i 

... 

200 

From  an  examination  of  the  table  it  is  an  easy  matter  to  pick  out 
the  smallest  weight  of  a  particular  element  occurring  in  the  mole- 
cule of  a  compound,  and  this  by  definition  is  the  atomic  weight  of 
the  element.  Thus,  among  the  compounds  of  carbon  examined 


DETERMINATION    OF   ATOMIC   WEIGHTS      117 

none  contain  less  than  12  parts  of  that  element  in  the  molecule, 
and  therefore  12  is  the  atomic  weight  of  carbon.  In  the  same 
way,  no  known  compound  contains  less  than  16  parts  of  oxygen 
in  the  molecule,  and  therefore  16  is  the  atomic  weight  of  oxygen. 
Further,  the  table  shows  that  the  atomic  weight  of  chlorine  is  35.5, 
and  that  of  mercury  200.  In  some  cases  a  molecule  contains  a 
simple  multiple  of  the  atomic  weight  of  an  element,  for  example, 
carbon  dioxide  contains  32  parts  of  oxygen.  In  this  case  we  make 
the  obvious  assumption  that  two  atoms  of  oxygen  are  present  in  the 
molecule. 

It  is,  of  course,  evident  that  this  method  gives  only  an  upper 
limit  for  the  atomic  weights,  but  if  sufficient  volatile  compounds 
are  examined,  some  are  certain  to  contain  only  one  atom  in  the 
molecule. 

Although  it  is  now  clear  that  atomic  weights  may  be  deduced  from 
the  results  of  vapour  density  determinations  alone,  by  application 
of  Avogadro's  hypothesis,  yet  far  greater  confidence  will  be  felt  in 
the  results  if  they  can  be  corroborated  by  other  methods.  Fortu- 
nately, there  are  at  least  four  other  methods  which  may  be  used  for 
this  purpose,  and  they  confirm  in  the  most  satisfactory  way  the 
conclusions  drawn  from  volumetric  data. 

The  more  important  methods  for  fixing  the  relative  values  of  the 
atomic  weights  are  as  follows  : — 

1.  Volumetric  methods  (already  dealt  with). 

2.  Purely  chemical  methods. 

3.  Methods  based  on  specific  heat  determinations. 

4.  Methods  based  on  isomorphism. 

5.  Methods  based  on  the  periodic  system  of  the  elements. 
Methods  (2),  (3)  and  (4)  will  now  be  briefly  considered  ;  the  fifth 
method  will  be  referred  to  at  a  later  stage  (p.  385). 

(2)  Chemical  Methods  of  Fixing  Atomic  Weights— 
The  first  step  in  using  this  method  is  exactly  the  same  as  for  the 
volumetric  method,  that  is,  an  analysis  is  made  of  the  substance  in 
order  to  determine  the  relative  amounts  of  the  components.  The 
further  procedure  may  be  illustrated  by  reference  to  water.  Water 
contains  in  round  numbers  i  part  by  weight  of  hydrogen  to  8  parts 
by  weight  of  oxygen.  If  the  formula  for  water  is  HO,  it  is  clear  that 
the  atomic  weight  of  oxygen  is  8,  referred  to  hydrogen  as  unit.  If, 
on  the  other  hand,  the  formula  is  H2O,  the  ratio  of  the  amounts  of 
the  two  elements  is,  as  before,  i  :  8  or  2  :  16,  and  therefore  the  atomic 
weight  of  oxygen  is  16.  If,  on  the  other  hand,  the  formula  is  H3O, 


n8     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  atomic  weight  of  oxygen  must  be  24  in  order  to  maintain  the 
experimentally  found  ratio  i  :  8.  Now  it  has  been  found  that  the 
hydrogen  can  be  displaced  from  water  in  two  stages  (by  means  of 
metallic  sodium,  for  instance),  and  in  two  stages  only.  As  not  less 
than  one  atom  can  be  displaced,  it  follows  that  there  are  at  least  two 
(and  probably  only  two)  atoms  of  hydrogen  in  the  molecule  of  water, 
the  formula  is  therefore  H2O,  and  the  atomic  weight  of  oxygen  is 
16.  The  last  conclusion  is  confirmed  by  the  experimental  fact  that 
the  oxygen  cannot  be  displaced  in  stages,  and  therefore  only  one 
atom  is  present. 

The  method  will  now  be  further  illustrated  by  application  to  two 
compounds  of  carbon  with  hydrogen,  called  methane  and  ethylene 
respectively.  This  case  is  of  historical  interest,  as  it  led  Dalton  to  the 
discovery  of  the  law  of  multiple  proportions.  Methane  contains 
3  parts  of  carbon  to  i  of  hydrogen  (in  other  words,  the  equiva- 
lent of  carbon  is  3),  and  ethylene  contains  6  parts  of  carbon  to 
i  of  hydrogen.  Dalton,  who  based  his  atomic  weights  to  some 
extent  on  assumed  simplicity  of  composition,  was  of  opinion  that 
ethylene  contained  one  atom  of  each  element,  and  hence  that  the 
atomic  weight  of  carbon  was  6.  We  now  know,  however,  that  one- 
fourth  of  the  hydrogen  in  methane  can  be  displaced  by  chlorine ; 
hence  methane  contains  four  hydrogen  atoms,  and  to  retain  the 
experimental  ratio  of  3  :  i  the  atomic  weight  of  carbon  must  be  12, 
provided  only  one  atom  is  present.  The  formula  for  methane  is 
therefore  CH4,  and  it  can  be  shown  by  similar  reasoning  that  the 
formula  of  ethylene  is  C2H4.  The  same  conclusion  can  be  reached 
more  readily  by  the  volumetric  method,  as  shown  in  the  table,  p.  116. 

(3)  Fixing  of  Atomic  Weights  from  Specific  Heats— 
Dulong  and  Petit,  in  1818,  made  the  remarkable  observation  that 
when  the  specific  heats  of  the  elements  in  the  solid  form  were 
multiplied  by  the  respective  atomic  weights,  the  products  were  in  all 
cases  approximately  the  same,  the  average  value  being  about  6.2. 
This  is  illustrated  in  the  table  on  the  following  page,  the  values  given 
for  the  specific  heats  being  the  mean  between  17°  and  100°. 

The  mean  value  of  the  product  for  the  metals  quoted  in  the  table  is 
about  6.1,  and  it  will  be  seen  that  only  in  the  case  of  silicon  is  there 
a  serious  deviation  from  the  average  value,  the  product  being  much 
smaller.  The  same  applies  to  three  other  light  elements,  the  names 
and  products  of  which  are  as  follows:  beryllium,  (9.1)  3.7;  boron, 
(u)  2.6;  carbon,  (12)  2-2.8.  It  has  been  shown,  however,  that  the 
specific  heats  of  these  elements  increase  very  rapidly  with  the  tern- 


DETERMINATION    OF   ATOMIC   WEIGHTS      119 


perature,  so  -that  the  deviation  from  the  normal  value  of  the  product 
is  much  smaller  at  high  temperatures. 


Element. 

Atomic 
Weight. 

Specific 
Heat. 

Product. 

Lithium     

Magnesium                  .         »         . 

7 
24.4 
27.1 

0.94 
0.247 
0.217 

6.6 
6.0 

5.Q 

28.4 

0.175 

5-° 

56 

O.IIO 

6.2 

Copper      
Silver          

,34 

0.093 
0.056 

i'8 

6.1 

Tin    . 
Antimony 
Platinum 
Gold. 
Mercury 
Bismuth 

119 

120 

J95 
197 

200 
208.5 

0.0556 
0.0503 
0.0310 
0.0310 

0.0335 
0.0303 

6.6 
6.0 
6.1 
6.1 

^7 

6-3 

Another  way  of  regarding  Dulong  and  Petit's  law  is  that  quantities 
of  different  substances  in  the  ratio  of  their  atomic  weights  require 
the  same  expenditure  of  heat  to  raise  them  through  an  equal  number 
of  degrees.  The  product  of  specific  heat  and  atomic  weight  may  be 
called  the  atomic  heat,  and  the  purport  of  Dulong  and  Petit's  law  is 
that  the  atomic  heats  of  all  solid  elements  are  approximately  equal 
(6.2  calories  being  required  to  raise  the  atomic  weight  in  grams  i° 
in  temperature).  It  is  certainly  a  most  remarkable  fact  that  the 
amounts  of  heat  required  to  raise  7  grams  of  lithium  and  208 
grams  of  bismuth  through  the  same  number  of  degrees  are  equal. 

The  method  of  using  Dulong  and  Petit's  law  to  fix  the  atomic 
weights  will  be  obvious  from  the  above.  Once  the  law  gained  the 
confidence  of  chemists,  it  proved  very  valuable  in  deciding  whether  a 
number  or  some  multiple  or  sub-multiple  represented  the  true  atomic 
weight. 

(4)  Fixing  of  Atomic  Weights  by  Isomorphism— It  was  ob- 
served by  Mitscherlich  (1822)  that  the  salts  known  as  sodium  hydrogen 
phosphate  and  sodium  hydrogen  arsenate  separate  from  solution  with 
the  same  number  of  molecules  of  water,  I2H2O,  are  identical  or 
nearly  so  in  crystalline  form,  and  when  a  solution  containing  the 
mixed  salts  is  evaporated  crystals  containing  both  substances  separate. 
It  is  clear  that  there  is  a  very  great  similarity  between  the  crystalline 
forms  of  the  two  salts,  which  is  expressed  by  saying  that  the  salts  are 
isomorphous,  or  have  the  same  crystalline  form.  On  the  basis  of 


i2o     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

these  and  similar  observations  Mitscherlich  brought  forward  his  Law 
of  Isomorphism,  according  to  which  compounds  of  the  same  crystal- 
line form  are  of  analogous  constitution.  Thus  when  one  element 
replaces  another  in  a  compound  without  alteration  of  the  crystalline 
form,  it  is  assumed  that  they  displace  each  other  atom  for  atom. 

A  little  consideration  shows  that  a  method  of  fixing  atomic  weights 
may  be  based  on  the  above  statements ;  if  an  element  of  known 
atomic  weight  is  displaced  by  another  of  unknown  atomic  weight,  the 
amounts  are  in  the  ratio  of  the  atomic  weights.  To  take  a  simple  case, 
the  two  salts,  potassium  chloride  and  potassium  iodide,  are  isomor- 
phous,  and  we  assume  that  the  atomic  weight  of  chlorine  is  known. 
Analysis  shows  that  potassium  chloride  contains  39  parts  of  potassium 
to  35.5  parts  of  chlorine,  and  that  for  the  same  amount  of  potassium 
the  other  salt  contains  127  parts  of  iodine.  As  iodine  is  assumed  to 
have  displaced  chlorine  atom  for  atom,  it  follows  that  127  is  the 
atomic  weight  of  iodine. 

Further  reference  to  crystal  form  and  isomorphism  is  made  at  a 
later  stage.  The  accepted  atomic  weights  for  the  elements  are  given 
in  the  table  published  by  an  International  Committee  (1914),  which 
for  convenience  of  reference  is  printed  on  the  back  page  of  the  cover 
of  the  book. 

Combining  Weights,  Chemical  Equivalents,  Atomic 
Weights — From  the  foregoing  sections  the  advantage  of  using 
atomic  weights  in  place  of  combining  weights  (chemical  equivalents) 
will  be  evident.  If,  for  example,  we  use  the  combining  weight  for 
oxygen,  8,  the  least  number  of  combining  weights  in  a  molecule  will 
be  2,  and  this  introduces  a  needless  complication.  In  the  same  way, 
each  molecule  of  a  carbon  compound  would  contain  4  or  some  multiple 
of  four  combining  weights. 

Still  another  advantage  of  the  atomic  weights  over  the  combining 
weights  is  seen  in  Dulong  and  Petit's  law.  There  is  no  such  con- 
nexion between  the  specific  heats  and  the  combining  weights  as  has 
been  shown  to  hold  for  specific  heats  and  atomic  weights.  Further, 
the  remarkable  relationships  between  the  atomic  weights  according 
to  the  periodic  system  (p.  379)  are  entirely  absent  when  combining 
weights  are  used  instead.  Finally,  and  in  some  respects  most  im- 
portant of  all,  each  element  has  only  one  atomic  weight,  whereas 
many  elements  have  more  than  one  equivalent. 

Several  methods  of  determining  chemical  equivalents  have  already 
been  indicated.  Most  of  them  are  based  on  the  definition  of  the 
chemical  equivalent  (combining  weight)  as  that  weight  of  an  element 


COMBINING   WRIGHTS  121 

which  combines  with  or  displaces  8  parts  by  weight  of  oxygeji  or 
1.008  parts  by  weight  of  hydrogen.  Thus  in  the  case  of  metals  which 
liberate  hydrogen  from  acids,  the  volume  of  the  hydrogen  set  free 
by  a  certain  weight  of  the  metal  may  be  measured,  or,  in  the  case 
of  substances  forming  stable  compounds  with  oxygen  the  oxides  may 
be  analysed.  Another  method  of  determining  chemical  equivalents 
is  based  on  the  observation  that  one  metal  may  displace  another 
from  one  of  its  salts  without  any  evolution  of  gas  or  alteration  of  the 
neutrality  of  the  solution.  The  amounts  of  the  metals  concerned 
in  the  change  are  clearly  in  the  ratio  of  the  chemical  equivalents, 
as  they  are  the  amounts  which  enter  into  combination  with  the  same 
quantity  of  the  negative  component  of  the  salt.  These  methods  are 
fully  described  in  a  later  part  of  the  chapter. 

It  has  already  been  stated  that  many  elements  may  be  regarded  as  having  more 
than  one  equivalent  or  combining  weight,  and  this  is  particularly  the  case  when 
the  compounds  of  the  element  in  question  with  oxygen  are  examined.  Thus 
there  are  five  oxides  of  nitrogen  containing  respectively  2.8,  3.5,  4.67,  7.0,  and 
14.0  parts  of  this  element  in  combination  with  8  parts  of  oxygen,  so  that  nitrogen 
has  at  least  5  chemical  equivalents.  We  might  choose  one  of  these  as  the 
equivalent  and  regard  the  others  as  multiples  or  submultiples  of  this  equivalent 
(cf.  p.  104),  or,  on  the  other  hand,  we  might  assume  that  this  and  other  elements 
have  more  than  one  equivalent.  The  latter  method  is  in  some  respects  preferable. 

Establishment  of  Molecular  Formulae— When  the  atomic 
weights  are  known,  it  is  a  relatively  simple  matter  to  establish  the 
molecular  formula  of  a  chemical  compound.  In  the  first  instance,  the 
percentage  composition  must  be  known.  This  can  be  determined  by 
direct  analysis  or  by  synthesis,  that  is,  by  finding  the  relative  amounts 
of  the  components  which  unite  to  form  the  compound.  The  next  step 
is  to  divide  the  relative  proportions  of  the  different  elements  present 
by  the  respective  atomic  weights  ;  this,  reduced  to  its  simplest  terms, 
gives  us  the  ratio  between  the  number  of  atoms  in  the  molecule. 
The  formula  thus  obtained  is  termed  the  empirical  formula,  and  the 
molecular  formula  may  be  the  same  as,  or  a  simple  multiple  of,  the 
empirical  formula.  In  order  to  settle  which  multiple  is  to  be  used, 
the  molecular  weight  has  finally  to  be  determined. 

The  method  will  be  fully  understood  from  an  example.  It  was  found  by 
analysis  that  hydrogen  peroxide  (p.  138)  contains  5.93  per  cent,  of  hydrogen  and 
94.07  per  cent,  of  oxygen.  Dividing  these  numbers  by  the  respective  atomic 
weights  in  order  to  obtain  the  relative  number  of  atoms,  we  have  5.88:5.88  or 
i:  i,  that  is,  the  same  number  of  atoms  of  each  element  are  present,  and  the 
empirical  formula  is  HO.  The  molecular  formula  is  therefore  (HO)*  where  x  is 


T22     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

an  integral  number  which  can  be  found  from  the  molecular  weight.  The  latter 
determined  directly  is  34,  which  is  satisfied  by  making  x=2,  so  that  the  formula 
of  hydrogen  peroxide  is 


When  the  components  and  the  substance  itself  can  be  obtained  in 
gaseous  form,  the  establishment  of  the  molecular  formula  is  still 
simpler.  It  has  been  found,  for  instance,  that  two  volumes  of  am- 
monia yield  on  decomposition  one  volume  of  nitrogen  and  three  volumes 
of  hydrogen  when  all  the  gases  are  measured  under  the  same  con- 
ditions. As  each  molcule  occupies  unit  volume,  we  have 

i  vol.  nitrogen     +     3  vols.  hydrogen      =     2  vols.  ammonia 


i  molecule  nitrogen  +  3  molecules  hydrogen  =  2  molecules  ammonia 
that  is       N2  +  3H2  =  2NH3 

In  the  case  of  substances  whose  molecular  weights  cannot  be 
determined  by  any  known  method  the  simplest  formula  (that  is,  the 
empirical  formula)  is  used,  e.g.  KNO3,  H2SO4. 

Equations  and  Calculations—  As  we  are  now  familiar  with 
the  methods  of  representing  chemical  compounds  by  their  formulae, 
we  are  in  a  position  to  represent  the  chemical  changes  considered  in 
the  previous  chapters  much  more  concisely.  The  combination  of 
hydrogen  and  chlorine  to  form  hydrogen  chloride,  for  example,  is 
represented  by  the  equation 

H2  +  C12  =  2HC1, 

and  the  combination   of  hydrogen   and   oxygen   to  form   water  as 
follows  :  — 


As  chemical  reactions  take  place  between  molecules,  it  is  preferable 
to  write  the  equation  representing  the  formation  of  water  as  above* 
instead  of  the  method  occasionally  used,  H2  +  O  =  H2O. 

In  representing  a  chemical  change  by  an  equation,  all  the  reacting 
substances  and  all  the  products  must  be  known,  as  well  as  their 
respective  molecular  formufee.  The  first  step  is  to  put  on  the  left- 
hand  side  the  formulas  for  the  reacting  substances,  and  on  the  right- 
hand  side  the  formulae  for  the  products.  The  next  step  is  to  satisfy 
the  law  of  the  conservation  of  mass,  that  the  sum  of  the  weights  of 
the  products  is  equal  to  the  sum  of  the  weights  of  the  reacting  sub- 
stances. The  simplest  way  of  doing  this  is  to  ensure  that  the 


FORMULAE   AND    EQUATIONS  123 

same  atoms  and  the  same  number  of  each  appear  on  the  two  sides  of 
the  equation. 

These  statements  will  now  be  illustrated  by  writing  the  equation 
representing  the  action  of  sulphuric  acid,  H2SO4,  on  zinc,  Zn,  giving 
zinc  sulphate,  ZnSO4,  and  hydrogen,  H2,  as  follows  :  — 

Zn  +  H2SO4=ZnSO4  +  H2. 

It  is  clear  that  each  atom  occurs  the  same  number  of  times  on  the  two 
sides  of  the  equation,  so  that  the  latter  is  complete  as  it  stands. 

A  slightly  more  complicated  case  is  the  effect  of  heat  on  potassium 
chlorate,  KC1O3,  giving  rise  to  potassium  chloride,  KC1,  and  oxygen, 
O2.  The  first  stage  is  as  follows  :  — 

KC1O3  =  KC1  +  O2, 

but  it  is  evident  that  in  this  form  the  atoms  do  not  occur  the  same 
number  of  times  on  each  side  ;  according  to  the  usual  expression  the 
equation  is  not  "balanced."  For  this  purpose,  sufficient  potassium 
chlorate  must  be  taken  to  yield  an  even  number  of  oxygen  atoms. 
The  completed  equation  is  as  follows  :  — 

2KC1O3=2KC1  +  3O2. 

Proceeding  in  a  similar  way,  the  reversible  reaction  between  water 
and  chlorine  at  high  temperatures  is  represented  by  the  equation 


and  the  action  of  steam  on  iron,  which  is  also  reversible,  by  the 
equation 


The  student  should  now  try  to  write  a  number  of  equations  for  himself, 
bearing  in  mind  that  all  the  reacting  substances  and  products,  as  well 
as  their  formulae  must  first  be  known. 

Calculations—  It  has  already  been  pointed  out  that  the  symbols 
in  a  chemical  formula  have  a  quantitative  significance,  inasmuch  as 
they  stand  for  quantities  of  the  different  substances  proportional 
to  their  atomic  weights.  This  consideration  at  once  enables  us  to 
calculate  from  a  chemical  equation  the  relative  amounts  of  the  react- 
ing substances  and  the  products.  This  may  be  illustrated  by  the 
equation  representing  the  action  of  sulphuric  acid  on  zinc. 

Zn     +     H2SO4      =      ZnSO4     +     H2 

2 

2 


i24     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

The  sum  of  the  weights  of  the  atoms  in  each  molecule  is  found,  and  the 
equation  shows  that  65  parts  of  zinc  react  with  98  parts  of  pure  sul- 
phuric acid  to  give  161  parts  of  zinc  sulphate  and  2  parts  of  hydrogen. 

When  gases  take  part  in  a  chemical  change,  we  can  further 
calculate  the  relative  volumes  concerned,  bearing  in  mind  the  result 
already  given  (p.  no)  that  the  molecular  weight  of  any  gas  in  grams 
occupies  22.4  litres  at  normal  temperature  and  pressure.  If,  in  the 
above  equation,  the  amounts  are  expressed  in  grams,  it  is  evident  that 
65  parts  of  zinc  liberate  2  grams  or  22.4  litres  of  hydrogen  measured 
under  normal  conditions. 

The  statements  will  now  be  illustrated  by  means  of  further 
examples,  (i)  What  weight  and  what  volume  of  oxygen,  measured 
under  normal  conditions,  will  be  obtained  by  completely  decomposing 
10  grams  of  mercuric  oxide  ?  The  equation  is  as  follows  :  — 


400  +  32  32 

and  shows  that  432   grams   of  mercuric  oxide   yield  32  grams  of 
oxygen,   so  that   10  grams  yield    .—  x  32  =0.74  grams  of  oxygen. 

Further,  since   432   grams  of  mercuric  oxide  yield  32  grams  =  22.4 

10 
litres  of  oxygen  at  N.T.P.,  10  grams  will  yield  ~-  x  22.4  =  0.518  litre 

or  518  c.c.  of  oxygen. 

(2)  What  weight  of  chlorine  will  be  obtained  by  heating  100  c.c.  of 
a  10  per  cent,  solution  of  hydrochloric  acid  with  excess  of  manganese 
dioxide,  and  what  weight  of  sodium  chloride  can  be  obtained  by 
complete  combination  of  the  chlorine  thus  formed  with  sodium  ? 

The  first  reaction  (compare  p.  88)  is  represented  by  the  equation 


4x36.5  2x35.5 

As  only  the  relationship  between  the  weights  of  hydrochloric  acid  and 
of  chlorine  is  required,  the  molecular  weights  of  the  other  substances 
need  not  be  calculated. 

As  146  grams  of  pure  hydrogen  chloride  yield  71  grams  of  free 
chlorine  under  the  conditions  described,  10  grams  of  hydrogen 
chloride  (the  amount  present  in  100  c.c.  of  a  10  per  cent,  solution) 

yield—  >x  10  —  4.86  grams  of  chlorine, 


FORMULAS   AND    EQUATIONS  125 

The  second  equation  concerned  is 

2Na  +  Cl2  =  2NaCl 

46      71         117 

And  as  71  grams  of  chlorine  yield  117  grams  of  sodium  chloride, 
4.86  grams  of  chlorine  yield  4.86  x  —  =  8.0  grams  of  sodium  chloride. 

The  relationships  between  the  volumes  of  gases  concerned  in 
chemical  changes  can  be  deduced  very  simply  on  the  basis  of  the 
deduction  from  Avogadro's  hypothesis  already  mentioned  (p.  109), 
that  the  molecule  of  any  substance  in  the  form  of  vapour  occupies 
unit  volume.  As  already  indicated,  this  has  no  reference  to  actual 
numerical  values,  but  is  merely  used  for  purposes  of  comparison.  On 
this  basis  the  volumes  of  the  gases  concerned  in  the  reversible  reaction 
between  steam  and  chlorine  (p.  86)  are  as  follows  : — 

2C12  +  2H2O^4HC1  +    O2 
2  vols.     2  vols.     4  vols.     i  vol. 

This  equation  shows  that  if  the  reaction  proceeds  completely  from 
left  to  right  there  is  an  expansion  of  4  volumes  to  5  ;  if  from  right  to 
left  there  is  a  contraction  from  5  volumes  to  4. 

In  the  foregoing  we  have  assumed  that  the  formulas  are  known,  and 
have  deduced  the  corresponding  volume  changes.  It  is  evident, 
however,  that  the  rule  may  be  employed  in  the  converse  way  to 
deduce  chemical  formulas  on  the  basis  of  the  observed  changes  of 
volume  accompanying  chemical  changes.  It  is  on  this  principle  that 
the  formula  for  ammonia  has  already  been  deduced  (p.  122),  and 
numerous  other  illustrations  will  be  met  with  later  (cf.  nitric  oxide} 
p.  232,  and  sulphur  dioxide,  p.  301). 

In  the  present  section  we  have  on  several  occasions  made  use  of 
the  molecular  weight  of  a  substance  in  grams.  As  these  amounts  are 
used  very  largely  in  chemistry,  it  is  convenient  to  have  a  shorter 
name,  and  for  this  purpose  the  term  mol>  introduced  by  Ostwald,  is 
suitable.  According  to  the  molecular  theory,  mols  of  different 
substances  contain  the  same  number  of  molecules  in  each  case,  and 
are  naturally  often  employed  for  comparative  purposes. 

Practical  Determination  of  Chemical  Equivalents— 
The  chemical  equivalent  of  an  element  has  already  been  defined  as 
that  amount  of  it  which  combines  with  or  displaces  8  parts  by  weight 
of  oxygen,  or  1.008  parts  by  weight  of  hydrogen.  It  is  desirable 


:26     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


that  the  student  should  determine  some  equivalents,  and  in  order  to 
illustrate  the  methods  employed,  a  few  examples  will  be  given. 

(i)  Determination  of  the  Equivalent  of  Zinc  —  The 
equivalent  of  zinc  (and  of  other  metals  which  readily  dissolve  in  acids, 
giving  off  hydrogen)  may  be  determined  by  dissolving  a  known 
weight  of  zinc  in  excess  of  hydrochloric  or  sulphuric  acid  and 
measuring  the  hydrogen  evolved.  A  convenient  form  of  apparatus 
for  this  purpose  is  shown  in  Fig.  40.  The  wide  graduated  tube  A, 
open  at  the  lower  end,  is  provided  with  a  side  tube  connected  by 

rubber  tubing  to  a  glass  tube  pass- 
ing through  the  rubber  cork  closing 
the  wide-mouthed  bottle  B.  At  the 
commencement  of  the  experiment  a 
weighed  amount  of  zinc  is  placed  in 
B,  as  is  also  a  small  bottle  containing 
more  than  sufficient  acid  to  dissolve 
the  zinc.  The  cork  is  then  inserted, 
and  the  water  in  the  wide  graduated 
tube  brought  to  a  convenient  height 
by  means  of  the  short  rubber  tube  and 
clip  at  the  top  of  A  ;  the  latter  is  then 
closed  by  the  clip.  The  bottle  B  is 
now  inclined,  so  that  the  acid  comes 
in  contact  with  the  zinc  ;  the  hydrogen 
given  off  displaces  the  water  in  A. 
When  all  action  has  ceased,  the 
graduated  tube  is  arranged  so  that  the 
water  stands  at  the  same  level  outside 
and  inside,  the  volume  of  the  hydro- 
gen is  then  read  off,  and  the  tempera- 


FIG.  40. 


ture  of  the  water  in  the  jar  C  noted.  The  volume  of  the  gas  is  then 
corrected  to  normal  temperature  and  pressure  (p.  45),  and  the  weight 
of  the  gas  calculated  in  the  usual  way  (p.  124).  The  weight  of  zinc 
which  would  give  I  part  by  weight  of  hydrogen  is  by  definition  the 
chemical  equivalent. 

(2)  Determination  of  the  Equivalent  of  Copper  by 
Analysis  of  the  Oxide — The  equivalent  of  copper  can  readily 
be  determined  by  reducing  a  known  weight  of  the  oxide  to  copper  by 
means  of  hydrogen  and  weighing  the  resulting  copper.  The  arrange- 
ment of  the  apparatus  is  shown  in  Fig.  17.  The  bulb  tube  is 
weighed  empty,  and  then  again  after  a  small  quantity  of  dried  copper 


FORMULAE   AND   EQUATIONS  127 

oxide  (say,  I  gram)  has  been  placed  in  it.  The  bulb  is  heated  and 
dry  hydrogen  passed  over  it  till  reduction  is  complete,  as  shown  by 
the  colour ;  it  is  then  allowed  to  cool  in  the  stream  of  hydrogen  and 
subsequently  weighed.  The  loss  of  weight  represents  the  oxygen  with 
which  the  copper  was  combined  ;  the  amount  of  the  latter  is  the 
difference  between  the  final  weight  of  the  bulb  and  the  weight  of  the 
latter  when  empty.  From  the  results  the  chemical  equivalent — the 
amount  of  copper  combined  with  8  parts  by  weight  of  oxygen— can 
readily  be  calculated  (p.  55). 

The  same  experiment  has  already  been  referred  to  in  connexion 
with  the  composition  of  water  (p.  58). 

(3)  Equivalent  of  Copper  by  Displacement  with  Mag- 
nesium— A  weighed  amount  of  magnesium  ribbon,  say,  0.25  gram,  is 
put  into  excess  of  a  warm  solution  of  copper  sulphate.  After  some  time 
it  will  be  found  that  the  magnesium  has  completely  disappeared,  and 
a  red  precipitate  of  copper  is  obtained.  The  precipitate  is  separated 
by  filtration,  washed,  dried,  and  weighed.  It  will  be  found  that  0.25 
gram  of  magnesium  displaces  about  0.65  gram  of  copper.  The 
chemical  equivalent  of  magnesium  is  about  12.2,  so  that  I  equivalent  of 

magnesium  displaces   i2.2X-L-^=3i.7    grams    or    i    equivalent  of 

copper. 
Practical  Determination  of  Vapour  Densities—  In  the 

present  chapter  we  have  had  many  illustrations  of  the  great  part 
which  a  knowledge  of  vapour  densities  has  played  in  the  development 
of  chemical  theory.  The  principle  of  the  methods  employed  in 
vapour  density  determinations  is  extremely  simple.  The  volume 
occupied  by  a  known  weight  of  the  vapour  under  known  conditions  of 
temperature  and  pressure  is  found,  and  the  ratio  of  this  weight  to  the 
weight  of  an  equal  volume  of  hydrogen  under  the  same  conditions  is 
the  required  density.  Space  only  admits  of  a  brief  account  of  two  of 
the  methods  in  actual  use  ;  for  a  full  account  of  vapour  density  deter- 
minations a  text-book  of  physics  should  be  consulted. 

(i)  Victor  Meyers  Method — The  method  most  largely  used  was 
introduced  by  Victor  Meyer  in  1878.  The  most  remarkable  feature  of 
the  method  is  that  it  is  not  the  volume  of  the  vapour  itself  which  is 
measured,  but  that  of  an  equal  volume  of  air  displaced  by  the  vapour. 

The  apparatus  consists  of  a  cylindrical  vessel,  A,  of  about  200  c.c. 
capacity,  ending  in  a  long  neck  provided  with  two  side  tubes,  as 
shown  in  Fig.  41.  One  of  these  side  tubes,  /,  from  which  the 
displaced  air  issues  during  an  experiment,  is  bent  in  such  a  way  that 


128     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


its  free  end  can  be  brought  under  the  surface  of  water  in  a  suitable 
vessel.  On  the  other  tube,  /2,  ls  fitted  a  rubber  tube  carrying 
a  glass  rod,  which  can  be  moved  outwards  and  inwards,  and,  at  the 
commencement  of  the  experiment,  serves  to  retain  in  place  the  small 
glass  bulb  shown  in  the  figure,  containing  a  weighed  quantity  of  the 
liquid  the  vapour  density  of  which  is  to  be  determined.  The  top  of 
the  main  tube  is  closed  by  a  cork,  which  is  kept  in  place  throughout 
an  experiment,  and  a  little  asbestos  or  mercury  is  placed  in  the 
bottom  to  guard  against  fracture  of  the  glass  when  the  bulb  drops. 
The  apparatus  is  kept  at  a  constant  temperature  throughout  the 
greater  part  of  its  length  by  means  of  the 
vapour  of  a  liquid  boiling  in  the  outer  bulb- 
tube  B  ;  the  temperature  should  be  at  least  20° 
above  the  boiling-point  of  the  liquid  to  be 
vaporized. 

At  the  commencement  of  an  experiment  the 
bulb  and  rod  are  placed  in  position  and  the 
cork  inserted;  the  jacketing  liquid  is  then  boiled 
till  air  ceases  to  issue  from  the  end  of  the  tube 
/  and  bubble  through  the  water,  showing  that  the  tempera- 
ture inside  the  bulb  A  is  constant.     A  graduated  measur- 
ing tube,  C,  full  of  water,  is  then  inverted  over  the  end  of 
the  delivery  tube,  and  the  small  bulb  allowed  to  drop  by 
drawing  back  the  glass  rod.     When  air  ceases  to  issue 
from  the  end  of  the  delivery  tube,  the  graduated  tube  is 
closed  by  the  thumb,  removed  to  a  deep  vessel  containing 
water,  allowed  to  stand  till  the  temperature  is  constant, 
and  the  volume  of  air  read  off  when  the  water  outside  and 
inside  is  at  the  same  level. 

The  temperature  inside  the  tube  A  is  the  same  before 
and  after  the  experiment ;  the  only  difference  in  the  conditions  is  that 
a  certain  volume  of  air  is  displaced  by  an  equal  volume  of  vapour. 
The  observed  volume  of  air  is  therefore  that  which  the  vapour  would 
occupy  after  reduction  to  the  temperature  and  pressure  at  which  the 
air  is  measured  (provided  that  the  vapour  and  air  are  equally  affected 
by  changes  of  temperature  and  pressure,  which  is  approximately  the 
case  under  suitable  conditions).  The  temperature  is  that  of  the 
water,  and  the  pressure  that  of  the  atmosphere  less  the  vapour 
pressure  of  water  at  the  temperature  of  observation.  It  is  evident 
that  it  is  not  necessary  to  know  the  temperature  at  which  the  sub- 
stance is  vaporized,  and  this  is  one  of  the  advantages  of  the  method. 


FIG.  41. 


VALENCY  129 

The  mode  of  calculating  vapour  densities  and  molecular  weights 
from  the  observed  data  may  be  illustrated  by  the  following  example  : 
0.220  grams  of  chloroform  when  vaporized  displaced  45.0  c.c.  of  air, 
measured  at  20°  and  755  mm.  pressure.  As  the  vapour  pressure  of 
water  at  20°  is  17.4  mm.  (p.  65),  the  actual  pressure  exerted  by  the 
gas  is  737.6  mm.  Therefore,  as  0.220  grams  of  vapour,  at  20°  and 
737.6  mm.  pressure,  occupy  45.0  c.c.,  we  have  to  find  what  weight  in 
grams  will  occupy  22,400  c.c.  at  273°  abs.  and  760  mm.  pressure,  and 
this  will  be  the  required  molecular  weight.  Now  45.0  c.c.  at  20°  and 

737.6mm.  pressure  measured  45  x^x^6o~=4o-7  c-c-  at  N.T.P., 
and  weigh  0.220  grams.  Therefore,  22,400  c.c.  at  N.T.P.  weigh 
0.220  x  '  =121,  which  is  the  required  molecular  weight,  as 

compared  with  the  value  (119.5)  calculated  from  the  formula  CHC18. 

The  vapour  density  can  be  calculated  from  the  experimental  data  as 
follows.  We  have  seen  that  40.7  c.c.  of  the  vapour  weigh  0.220  grams 
at  N.T.P.  As  2.016  grams  of  hydrogen  measure  22,400  c.c.  at  N.T.P., 

40.7  c.c.  of  hydrogen  weigh  2.016  x  =0.00366  grams.     Hence 

22,400 

Wt.  of  substance  0.220 

Vapour  density = ^ — ? i — i — FTT  = Z2  =  o°- J  • 

3     Wt.  of  equal  vol.  of  H2    0.00366 

The  molecular  weight,  being  by  definition  double  the  vapour  density, 
is  120,  as  has  just  been  shown  by  the  alternative  method.1 

(2)  Hofmanris  Method — A  graduated  glass  tube  is  filled  with 
mercury,  and  inverted  in  a  bath  of  the  same  metal.  A  small  glass 
tube,  containing  a  weighed  quantity  of  the  substance  whose  vapour 
density  has  to  be  determined,  is  inserted  under  the  lower  edge  of  the 
graduated  tube ;  it  rises  to  the  top,  and  the  liquid  vaporizes,  displacing 
part  of  the  mercury.  The  volume  of  the  vapour,  its  temperature  and 
pressure,  are  then  read  off,  and  from  these  data  and  the  weight  of  the 
substance  the  vapour  density  and  molecular  weight  are  calculated 
in  the  usual  way.  The  advantage  of  Hofmann's  method  is  that  it 
admits  of  the  determination  of  densities  under  reduced  pressure. 

Valency  and  Structural  Formulae  — We  have  learnt  in 
previous  chapters  that  the  formula  of  hydrogen  chloride  is  HC1, 
and  of  water  H2O.  One  atom  of  chlorine  always  combines  with  one, 
and  not  more  "than  one,  atom  of  hydrogen  ;  similarly,  one  atom  of 
oxygen  combines  with  two  atoms  of  hydrogen,  neither  more  nor  less. 

1  The  slight  difference  between  the  two  results  is  due  to  the  fact  that  densities 
are  referred  to  that  of  hydrogen  as  unit,  molecular  weights  to  the  atom  of  oxygen 
as  i6(H  =  1.008). 

9 


130     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

It  is  evident  that  chlorine  and  oxygen  have  a  different  combining 
capacity  for  hydrogen.  The  combining  capacity  or  chemical  value  of 
an  element  is  termed  its  valency,  and  is  usually  measured  with  regard 
to  hydrogen.  An  element  such  as  chlorine,  one  atom  of  which  com- 
bines with  one  atom  of  hydrogen,  is  said  to  be  univalent,  whilst  oxygen, 
one  atom  of  which  combines  with  two  atoms  of  hydrogen,  is  said  to  be 
bivalent.  Elements  of  valency  three  are  trivalent  or  tervalent,  of 
valency  four  quadrivalent ',  and  so  on. 

Many  elements,  however,  either  do  not  form  compounds  with 
hydrogen  (e.g.  bismuth,  lead),  or  form  compounds  of  unknown 
molecular  formula  (e.g.  sodium  potassium).  In  such  cases  the 
valency  can  usually  be  determined  from  an  examination  of  compounds 
of  the  type  MXn  (M  =  Element;  X  =  halogen)  as  the  halogens  are 
univalent  with  regard  to  hydrogen. 

From  these  considerations,  the  valency  of  an  element  is  measured 
by  the  number  of  hydrogen  or  other  univalent  atoms  or  groups  with 
which  one  atom  of  it  can  combine  to  form  a  molecule. 

The  question  now  arises  as  to  whether  the  valency  of  an  element  is 
constant  or  variable.  This  question  was  widely  debated  by  chemists 
for  many  years,  but  it  is  now  agreed  that  the  same  element  may  have 
different  valencies  in  different  compounds.  The  highest  valency 
shown  by  any  element  is  eight. 

Hydrogen  is  invariably  univalent,  as  are  sodium  and  potassium. 
Oxygen  is  almost  always  divalent.  Chlorine  is  univalent  with  regard 
to  hydrogen,  but  in  binary  compounds  with  oxygen  and  in  ternary 
compounds  with  oxygen  and  hydrogen,  appears'  to  have  several 
different  valencies.  Barium  and  calcium  appear  to  be  always 
bivalent.  The  valencies  of  the  elements  will  be  discussed  in  con- 
nexion with  their  detailed  consideration. 

A  knowledge  of  the  common  valencies  of  the  elements  is  of  great 
value  in  writing  chemical  formulae.  We  have  seen  that  salts  are 
derived  from  acids  by  displacement  of  the  acidic  hydrogen  by  metals. 
Suppose  we  wish  to  know  the  formula  of  barium  chloride.  It  is 
derived  from  hydrochloric  acid  by  displacing  the  hydrogen  by  barium. 
As  the  latter  is  a  divalent  element,  and  therefore  displaces  two  atoms 
of  hydrogen,  the  chloride  is  derived  from  2HC1,  and  its  formula  is 
consequently  BaCl2. 

Groups  of  elements  may  also  be  regarded  as  having  a  definite 
valency.  Thus  since  the  formula  of  sulphuric  acid  is  H2SO4,  the  group 
SO4  is  bivalent,  as  it  is  associated  with  two  atoms  of  hydrogen. 
Similarly,  the  group  C1O3  is  univalent,  as  it  is  associated  with  one 


VALENCY  131 

atom  of  potassium,  a  univalent  element,  in  potassium  chlorate, 
KC1O3. 

The  valencies  in  a  compound  may  be  shown  by  means  of  bars  or 
links  between  the  constituent  atoms,  each  bar  representing  a  single 
valency.  The  formula  for  water  is  written  thus :  H  -  O  -  H,  showing 
that  the  two  valencies  of  oxygen  are  satisfied  by  the  two  atoms  of 
hydrogen.  Such  formulse  are  termed  graphic,  structural  or  constitu- 
tional formulae.  The  graphic  formula  of  hydrogen  chloride  is  H  -  Cl. 

The  formulas  of  the  oxides  of  chlorine,  C12O  and  C1O2,  may  serve 
as  further  illustrations.  The  first  oxide  is  written  thus  :  Cl  -  O  —  Cl, 
showing  that  it  is  of  the  same  type  as  water.  The  other  oxide,  C1O.2, 
presents  more  difficulty  from  this  point  of  view.  It  may  be  written 
O  =  C1  =  O,  showing  the  chlorine  as  quadrivalent ;  but  other  methods 
of  formulation  may  be  used. 

We  shall  see  in  detail  later  that  the  graphic  formula  of  a  compound 
is  an  attempt  to  represent  in  a  brief  way  its  characteristic  behaviour. 
As  the  behaviour  of  chemical  compounds  is  many-sided,  it  is  not 
surprising  that  opinions  often  differ  with  regard  to  the  most  suitable 
graphic  formulas  of  particular  compounds. 

Chemical  Equivalent,  Atomic  Weight,  and  Valency— 
The  chemical  equivalent  of  an  element  has  been  defined  as  that 
quantity  of  it  which  combines  with  or  displaces  one  part  by  weight 
of  hydrogen.  The  atomic  weight  of  an  element,  on  the  other  hand, 
may  displace  one  or  more  parts  by  weight  (one  or  more  atoms)  of 
hydrogen,  according  as  the  element  is  univalent  or  polyvalent.  If 
univalent,  the  atomic  weight  of  an  element  displaces  one  part  by 
weight  of  hydrogen  : 

Na  +  HOH->NaOH  +  H 
if  divalent,  two  parts  by  weight  of  hydrogen  : 

Zn+ H2SO4->ZnSO4-t- H2 

and  so  on. 

From  the  foregoing  it  is  evident  that 

Atomic  weight 

Valency          Chemical  eclmvalent- 

This  important  relationship  should  be  carefully  remembered.  If  the 
atomic  weight  and  the  chemical  equivalent  of  an  element  have  been 
determined  independently  by  the  methods  described  in  the  present 
chapter,  the  valency  can  at  once  be  deduced.  If,  on  the  other  hand, 


132     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  atomic  weight  and  valency  are  known,  the  chemical  equivalent 
can  at  once  be  calculated. 

In  the  light  of  this  statement,  the  chemical  equivalents  on  p.  115 
should  be  compared  with  the  corresponding  atomic  weights. 

Summary  —The  last  two  chapters  deal  with  many  ol  the  funda- 
mental principles  of  chemistry,  and  it  will  be  useful  briefly  to 
summarize  some  important  points.  The  basis  of  chemical  theory 
is  Avogadro's  hypothesis,  which  states  that  equal  volumes  of  all 
gases,  under  the  same  conditions  of  temperature  and  pressure,  con- 
tain the  same  number  of  molecules.  This  principle  enables  us  at  once 
to  fix  the  molecular  weights  of  volatile  substances,  which  are  referred 
to  the  atom  of  hydrogen  as  unit.  From  the  data  on  molecular 
weights,  the  atomic  weights  of  the  elements  can  be  determined,  the 
atomic  weight  being  by  definition  the  smallest  quantity  of  an  element 
which  occurs  in  a  molecule,  referred  to  the  atom  of  hydrogen  as  unit. 
When  the  atomic  weights  are  known,  the  molecular  formula  of  a 
compound  can  readily  be  established  as  follows.  The  percentage 
composition  is  determined  by  analysis,  the  relative  proportion  by 
weight  of  each  element  is  then  divided  by  the  corresponding  atomic 
weight,  and  the  result,  reduced  to  its  simplest  terms,  gives  us  the 
ratio  between  the  number  of  atoms  in  the  molecule,  the  so-called 
empirical  formula.  The  molecular  formula  is  the  same  as,  or  an 
integral  multiple  of,  the  empirical  formula  (p.  121).  It  must  be 
emphasized  that  the  methods  of  determining  molecular  weights  only 
give  approximate  values,  which  are  sufficiently  accurate  to  show 
which  multiple  of  the  empirical  formula  is  to  be  taken.  Molecular 
formulae  can  also  be  established  by  applying  the  deduction  from 
Avogadro's  hypothesis  that  the  molecule  of  any  substance  in  the 
state  of  vapour  occupies  unit  volume. 

The  atomic  weights  deduced  on  the  basis  of  Avogadro's  hypothesis 
are  in  full  agreement  with  those  determined  by  the  alternative 
methods.  Similarly  the  molecular  formulae  obtained  as  above  are  in 
agreement  with  the  conception  that  chemical  reactions  take  place 
between  molecules.  It  is,  in  fact,  possible  to  determine  the  molecular 
weight  of  a  substance  from  its  chemical  behaviour,  on  the  assumption 
that  the  molecule  is  the  smallest  quantity  of  a  substance  which  can 
take  part  in  a  chemical  change. 

The  exact  connexion  between  atomic  weights  and  chemical 
equivalents,  and  the  advantages  of  representing  composition  in  terms 
of  atomic  weights  instead  of  in  terms  of  equivalents,  have  been  fully 
explained  in  the  course  of  the  chapter. 


CHAPTER  XI 

OZONE   AND   HYDROGEN    PEROXIDE- 
THERMOCHEMISTRY 

IN  the  present  chapter  we  are  concerned  with  two  interesting  sub- 
stances closely  allied  to  two  elements,  hydrogen  and  oxygen, 
already  dealt  with  in  detail.  We  shall  learn  that  ozone  is  simply  a 
form  of  oxygen,  whereas  hydrogen  peroxide,  as  its  name  indicates, 
is,  like  water,  a  chemical  compound  of  hydrogen  and  oxygen.  Both 
substances  are  energetic  oxidizing  agents. 

OZONE 

Formula,  O3.     Molecular  weight =48.     Density=24. 

History — It  has  long  been  known  that  when  an  electrical  machine 
is  in  operation  a  peculiar  smell,  somewhat  like  that  of  dilute  chlorine, 
is  noticeable  in  its  vicinity,  and  in  1785  Van  Marum  showed  that  it 
was  due  to  the  action  of  the  electrical  discharge  on  the  oxygen  of  the 
atmosphere.  The  same  smell  is  sometimes  noticed  after  a  thunder- 
storm. In  1840,  Schonbein  showed  that  the  smell  was  due  to  the 
formation  of  a  definite  substance,  which  he  named  ozone  (fi&tv,  to 
smell),  and  described  other  methods  by  which  the  new  substance 
could  be  obtained. 

Occurrence — It  is  generally  assumed  that  traces  of  ozone  are 
normally  present  in  the  atmosphere,  but  recent  investigations  render 
this  extremely  doubtful  (see  below). 

Preparation — (i)  Ozone,  mixed  with  a  large  excess  of  oxygen, 
can  be  obtained  by  passing  a  silent  electric  discharge  through  oxygen. 
A  suitable  arrangement  for  demonstration  purposes  is  represented  in 
Fig.  42.  It  consists  of  a  wide  glass  tube  AA,  into  which  a  cork 
carrying  a  narrower  tube  B  is  fitted  as  shown.  The  outer  tube  is 
covered  with  tinfoil,  and  connected  to  one  pole  of  an  induction  coil ; 
the  inner  tube  contains  some  conducting  material  which  is  connected 
to  the  other  pole  of  the  coil.  Under  these  circumstances  the  discharge 

133 


i34     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

takes  place  between  the  two  tubes,  not  in  sparks,  but  as  a  luminous 
glow  ;  this  is  termed  a  silent  discharge.  Oxygen  or  air  is  then 
passed  through  the  space  between  the  tubes,  and  it  can  be  shown 
by  the  characteristic  smell,  and  by  application  of  the  tests  described 
below,  that  the  issuing  gas  contains  ozone. 

The  oxygen  is  always  present  in  large  excess,  but  under  favourable 
conditions  a  mixture  containing  5  to  6  per  cent,  of  ozone  can  be  obtained 
by  means  of  the  silent  discharge.  The  yield  of  ozone  is  increased  by 
keeping  the  temperature  low. 

(2)  The  oxygen  obtained  at  the  anode  when  dilute  sulphuric  acid 
is  electrolyzed  contains  more  or  less  ozone,  depending  on  the  condi- 
tions.    It  has  recently  been  shown  by  Fischer  and  Massenez  that 
under  favourable  conditions  a  mixture  containing  no  less  than  28  per 
cent,  by  weight  of  ozone  can  be  obtained  in  this  way.    The  best  yields 
were  obtained  with  solutions  containing  10  to  15  per  cent,  by  weight 

of  the  acid  and  a  high 
current  density.  The 
anode  was  of  platinum, 
the  exposed  surface 
being  very  small  in 
order  to  minimize  sub- 
sequent decomposition 
of  the  ozone,  and  ar- 
rangements were  made 
FlG  4a  for  keeping  the  tem- 

perature low. 

(3)  The   oxygen  formed  by  a  number  of  the   chemical    methods 
already  mentioned,  for  example,  by  the  action  of  concentrated  sul- 
phuric acid  on  potassium  permanganate,  manganese  dioxide,  or  barium 
peroxide,  contains  sufficient  ozone  to  answer  the  ordinary  tests. 

(4)  Ozone  is  formed  during  the  slow  oxidation  of  phosphorus  at  the 
ordinary  temperature.      If  a  freshly  scraped  stick  of  phosphorus  is 
exposed  to  the  air  for  a  short  time  in  a  covered  gas  jar,  the  smell  of 
ozone   can   readily  be   detected,   and  a   strip   of  moist   filter-paper 
impregnated  with  potassium  iodide  and  starch  (see  tests  for  ozone) 
suspended  in  the  jar  will  turn  blue. 

Physical  Properties— At  ordinary  temperatures  ozone  is  a  gas 
which,  in  thick  layers,  is  bluish  in  colour.  Under  atmospheric  pres- 
sure it  boils  at  -119°,  so  that  it  can  be  obtained  in  the  liquid  form  by 
passing  the  gas,  mixed  with  oxygen,  through  a  U-tube  immersed  in 
liquid  oxygen  (boiling-point  -  182.5°).  In  tnis  wav  a  liquid  is  obtained 


OZONE   AND   HYDROGEN    PEROXIDE          135 

containing  only  a  small  proportion  of  liquid  oxygen,  and  the  latter 
can  be  almost  completely  removed  by  evaporation,  being  much  more 
volatile  than  ozone.  Liquid  ozone  is  deep  indigo-blue  in  colour,  and 
is  extremely  explosive.  Ozone  is  much  more  soluble  in  water  than  is 
oxygen. 

Chemical  Properties  —  Pure  gaseous  ozone  is  so  unstable  as 
to  be  practically  unknown,  but  ozone  mixed  with  excess  of  oxygen 
is  fairly  stable  at  ordinary  temperatures.  When  the  mixture  is 
heated  to  250-300°,  however,  it  is  rapidly  reconverted  to  oxygen. 
This  may  readily  be  demonstrated  by  passing  the  gas  from  the 
ozone  apparatus  through  a  glass  tube  heated  by  a  Bunsen  flame, 
when  the  issuing  gas  will  no  longer  give  the  tests  for  ozone. 

The  facts  just  mentioned  show  that  the  equation  representing  the 
formation  of  ozone  from  oxygen 


is  reversible  ;  the  action  represented  by  the  upper  arrow  is 
favoured  by  the  silent  electric  discharge  (which  supplies  the  neces- 
sary energy),  but  when  a  fairly  large  proportion  of  ozone  is  present 
the  reaction  proceeds  rapidly,  and  almost  completely,  in  the  direction 
indicated  by  the  lower  arrow  at  250°.  The  interesting  relationships 
between  oxygen  and  ozone  are  considered  more  fully  in  the  chapter 
on  chemical  equilibrium  (p.  174). 

When  ozone  splits  up  into  oxygen,  a  large  amount  of  heat  is  given 
out.  The  reaction  2O3->3O2  liberates  about  68,200  cal.  (Jahn)  when 
gram-molecular  quantities  are  used,  that  is,  when  96  grams  of  ozone  are 
transformed  to  oxygen.  It  follows  at  once  by  Le  Chatelier's  theorem 
(p.  171)  that  raising  the  temperature  must  favour  the  reaction  which 
proceeds  with  absorption  of  heat  ;  that  is,  increase  of  temperature 
must  favour  the  formation  of  ozone  from  oxygen,  and  this  conclusion 
is  fully  borne  out  by  experiment.  The  apparent  contradiction  between 
this  result  and  the  rapid  decomposition  of  ozone  at  250°  will  be  dealt 
with  later  (p.  174). 

The  most  characteristic  chemical  property  of  ozone  is  that  it  is  a 
powerful  oxidizing  agent,  being  much  more  active  than  free  oxygen 
in  this  respect.  It  oxidizes  both  mercury  and  silver  at  room  tempera- 
ture, the  latter  becoming  blackened  owing  to  formation  of  an  oxide, 
Ag2O2.  Ozone  readily  bleaches  indigo,  litmus,  and  some  other 
colouring  matters,  and  destroys  india-rubber  connexions.  When 
passed  into  a  solution  of  potassium  iodide  containing  a  little  mucilage 


136    A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

of  starch,   the  solution  turns  blue.     The   reaction   in  this  case  is 
represented  by  the  equation 


that  is,  potassium  hydroxide,  oxygen  and  free  iodine  are  formed, 
and  the  latter  gives  rise  to  the  deep  blue  colour  with  starch. 

The  reaction  just  described  is  often  made  use  of  as  a  test  for  ozone 
as  follows.  A  little  starch  is  boiled  with  a  few  c.c.  of  water  for 
some  minutes,  and  a  few  crystals  of  potassium  iodide  dissolved  in  the 
thick  liquid.  Strips  of  filter-paper  are  then  smeared  with  the  solution 
and  when  dried  constitute  the  so-called  iodide  of  potassium  starch 
paper,  which  at  once  turns  blue  when  moistened  and  exposed  to 
ozone.  We  shall  find,  however,  that  this  test  paper  is  turned  blue 
by  many  oxidizing  agents  besides  ozone. 

The  last  equation  serves  to  illustrate  the  important  fact  that 
when  ozone  acts  as  an  oxidizing  agent  oxygen  is  generally  set 
free.  Further,  since  both  O3  and  O2,  being  single  molecules,  occupy 
two  unit  volumes,  the  volume  of  the  resulting  oxygen  is  the  same 
as  that  of  the  ozone.  It  is  therefore  evident  that  the  oxidation  is 
effected  by  the  extra  atom  of  oxygen  in  the  ozone  molecule. 

The  oxidizing  power  of  ozone  is  closely  connected  with  the  fact 
that  it  has  much  more  energy  than  an  equal  amount  of  oxygen. 

On  the  commercial  scale  ozone  is  employed  for  freeing  water  from 
micro-organisms  (p.  61),  for  bleaching  flour  and  for  other  purposes. 

Formula  of  Ozone—  The  facts  that  ozone  can  be  obtained 
from  pure  oxygen  alone,  and  that  on  heating  oxygen  and  oxygen 
only  is  obtained,  prove  that  it  consists  simply  of  oxygen.  In  the 
foregoing  we  have  assumed  that  whereas  the  molecule  of  oxygen 
contains  two  atoms,  O2,  that  of  ozone  contains  three  atoms  of  oxygen, 
and  is  therefore  represented  by  O3.  It  remains  to  indicate  the 
evidence  on  which  this  formula  is  based. 

(i)  The  question  could,  of  course,  at  once  be  settled  by  determining 
the  density  of  the  pure  gas,  but,  as  has  been  pointed  out,  the  latter 
cannot  be  prepared.  The  difficulty  has,  however,  been  got  over  by 
determining  the  density  of  a  mixture  of  oxygen  and  ozone  con- 
taining a  known  proportion  of  ozone  (Ladenburg).  Suppose,  for 
example,  the  weight  of  a  litre  of  the  mixture  at  N.T.P.  containing 
loo  c.c.  of  ozone  is  found  to  be  1.500  grams.  The  weight  of  i 
litre  of  oxygen  at  N.T.P.  is  1.429  grams,  and  of  900  c.c.  1.286  grams. 
The  remainder,  1.500-1.286  =  0.214  grams,  represents  the  weight  of 
100  c.c.  of  ozone  at  N.T.P.  The  weight  of  ozone  which  would 


OZONE   AND   HYDROGEN   PEROXIDE          137 

occupy  22.4  litres  at  N.T.P.,  that  is,  the  molecular  weight  of  ozone 
(p.  1 10),  is  therefore  0.2 14X^42°.  =  48  grams  approximately,  whence 

it  follows  that  the  formula  for  ozone  is  O3. 

(2)  Another  proof  depends  upon  the  fact  that  ozone  is  completely 
absorbed  from  its  mixture  with  oxygen  by  shaking  with  ordinary 
turpentine.      Now,  it  has  been  found  that  when  ozone  is  formed 
from  oxygen  there  is  a  definite  diminution  of  volume,  and  when 
the  ozonized  oxygen  is  shaken  with  turpentine  a  further  diminution 
of  volume  double  the  former  one  is  observed.     Suppose,  for  example, 
100  c.c.  of  oxygen  become  95  c.c.  when  partly  ozonized  ;  the  volume 
is  reduced  to  85  c.c.  by  shaking  with  turpentine.     It  is  evident  that 
the  10  c.c.  of  ozone  absorbed  by  the  turpentine  were  formed  from 
100-85  =  15  c.c.  of  oxygen,  so  that  three  volumes  of  oxygen  yield 
two  of  ozone.     This  observation  can  only  be  satisfied  if  the  change 
of  oxygen  to  ozone  is  represented  by  the  formula  3O2->2O3,  ac- 
cording to   which   3  volumes  of  oxygen  yield  2  volumes  of  ozone 
(p.  109). 

(3)  It  has  been  pointed  out  (diffusion  of  gases,  p.  47)  that  the 
relative  rate  of  diffusion  of  gases  is  inversely  as  the  square  roots 
of  their  densities.     Soret  found  that  ozone  diffused  more  slowly  than 
oxygen,  and  that  the  relative  rates  were  in  accordance  with  the  con- 
clusion that  the  density  of  ozone  is  24  when  O  =  i6. 

Tests  for  Ozone — The  tests  for  ozone  are  rather  important  in 
connexion  with  the  disputed  question  as  to  the  presence  of  this  gas 
in  the  atmosphere.  The  oxidizing  properties  already  mentioned, 
such  as  the  liberation  of  iodine  from  potassium  iodide  solution,  the 
bleaching  of  indigo,  etc.,  are  not  characteristic,  as  one  or  other  of 
them  is  also  given  by  other  substances  such  as  hydrogen  peroxide 
and  oxides  of  nitrogen,  although  not  by  oxygen.  The  characteristic 
smell  is  the  most  sensitive  test,  but  the  following  chemical  actions 
are  useful : — 

(a)  When  a  strip  of  potassium  iodide  starch  paper,  coloured  violet 
with  litmus,  is  exposed  to  an  atmosphere  containing  ozone,  potassium 
hydroxide,  as  well  as  iodine,  is  formed  (see  above),  and  the  former 
turns  the  litmus  paper  blue.  Hydrogen  peroxide  also  answers 
this  test. 

(ti)  When  very  pure  mercury  is  exposed  to  ozone,  the  metal  be- 
comes less  mobile,  and  adheres  to  the  glass.  This  effect  is  doubtless 
due  to  the  formation  of  traces  of  oxide. 

(c)  A  bright  silver  surface,  previously  polished  and  heated  to  red- 


138     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

ness,  is  blackened  on  exposure  to  ozone,  owing  to  the  formation  of 
silver  peroxide.  This  test  is  not  very  sensitive. 

(d)  Filter-paper  dipped  in  an  alcoholic  solution  of  the  so-called  tetra- 
base  (tetramethyl  p-p'-diaminodiphenylmethane)  becomes  violet  in 
ozone,  yellow  in  nitric  oxide,  and  is  not  altered  by  hydrogen  peroxide. 

As  already  mentioned,  it  is  very  doubtful  whether  ozone  is  a 
normal  constituent  of  the  atmosphere.  It  is  now  known  that  the 
positive  results  of  the  earlier  tests  were  due  mainly  to  the  presence  of 
nitrous  fumes  and  of  hydrogen  peroxide. 


HYDROGEN  PEROXIDE 

Formula,  H2O2. 

History — Hydrogen  peroxide  was  discovered  by  Thenard  (1818), 
who  obtained  it  by  the  action  of  acids  on  barium  peroxide. 

Occurrence  —  Minute  traces  of  this  compound  appear  to  be 
present  under  normal  conditions  in  the  atmosphere. 

Preparation — (i)  Hydrogen  peroxide  is  formed  in  small  amount 
when  hydrogen  burns  in  the  air.  This  is  best  shown  by  directing 
the  flame  on  the  surface  of  cooled  water  or  on  to  a  piece  of  ice  ;  the 
rapid  cooling  thus  secured  prevents  the  splitting  up  into  water  and 
oxygen  which  takes  place  in  the  flame  under  ordinary  conditions. 

(2)  Hydrogen  peroxide  is  obtained  on  the  large  scale  by  the  action 
of  dilute  sulphuric  acid  on  barium  peroxide,  represented  by  the 
equation 

Ba02+H2S04=BaS04  +  H202. 

The  best  results  are  obtained  by  slowly  adding  powdered  hydrated 
barium  peroxide,  BaO2,8H2O,  to  a  mixture  of  I  part  of  concentrated 
sulphuric  acid  and  5  parts  of  water  till  the  acid  is  nearly  neutralized 
(see  p.  99),  the  temperature  being  kept  low  by  immersing  the  vessel 
in  cold  water  or  in  a  mixture  of  water  and  ice.  Barium  sulphate  is 
insoluble  in  water,  and  can  be  removed  by  filtration,  a  dilute 
aqueous  solution  of  hydrogen  peroxide  being  thus  obtained.  A 
more  concentrated  solution  of  the  peroxide  may  be  obtained  by 
distilling  the  dilute  solution  under  reduced  pressure  (p.  65) ;  at 
first  almost  pure  water  passes  over,  and  then  a  fairly  concentrated 
solution  of  the  peroxidft.  By  repeating  this  process  of  fractional 
distillation  several  times,  hydrogen  peroxide  can  be  obtained  quite 
free  from  water  (cf.  p.  140). 

Instead  of  sulphuric  acid,  carbon  dioxide  or  phosphoric  acid  might 


OZONE   AND   HYDROGEN   PEROXIDE          139 

be  used  to  decompose  barium  peroxide,  as  both  form  insoluble  barium 
salts  which  can  be  separated  by  filtration  : 


+H2O  +  CO2->BaCO3+H2O2, 
3BaO2+2H3PO4->Ba3(PO4)2+3H2O2. 

Hydrochloric  acid  might  also  be  used  for  the  same  purpose,  but  as 
barium  chloride  is  soluble  in  water,  the  products  could  not  readily  be 
separated.  Merck's  method,  used  on  the  commercial  scale,  is  to  add 
sodium  peroxide,  Na2O2  (p.  397),  to  cooled  20  per  cent,  sulphuric 
acid  ;  the  sodium  sulphate  separating  is  filtered  off  and  the  solution 
concentrated  by  fractional  distillation.  The  small  amount  of  sodium 
sulphate  remaining  in  the  solution  does  not  appreciably  decompose 
the  peroxide  during  concentration. 

(3)  Hydrogen  peroxide  is  formed  in  small  amount  when  zinc  is 
shaken  with  water  and  air  or  oxygen,  and  also  when  lead,  copper, 
and  some  other  metals  are  shaken  with  air  and  dilute  sulphuric  acid. 
These    remarkable    reactions    are    by   no    means   well    understood. 
Hydrogen  peroxide  can  be  detected  in  many  other  cases  of  oxida- 
tion by  free  oxygen. 

(4)  Hydrogen  peroxide  is  formed  in  fair  amount  when  oxygen  is 
bubbled  through  the  solution  surrounding  a  negative  pole  at  which 
hydrogen  is  being  liberated  by  electrolysis.      The   change  can  be 
represented  by  the  equation  H2+O2  =  H2O2,  but  is  not  fully  under- 
stood.    Under  ordinary  conditions  hydrogen  and  oxygen  do  not  unite 
to  form  hydrogen  peroxide,  and  we  must  assume  that  the  hydrogen 
in  the  act  of  liberation  is  endowed  with  exceptional  activity. 

Physical  Properties  —  When  quite  free  from  water,  hydrogen 
peroxide  is  a  colourless  syrupy  liquid  of  specific  gravity  1.458  at  o°. 
Under  29  mm.  pressure  it  boils  at  65°  and  under  65  mm.  at  85°.  It 
can  be  obtained  in  colourless  crystals  by  cooling  the  pare  liquid  in  a 
mixture  of  ether  and  solid  carbon  dioxide  (p.  337)  ;  the  crystals  melt 
at  -  2°.  The  heat  of  solution  in  water  is  460  cal.  per  mol.  The  heat 
given  out  when  hydrogen  peroxide  splits  up  into  water  and  oxygen 
(see  below)  is,  according  to  Berthelot,  21,700  cal.  per  mol  (34  grains). 

Chemical  Properties—  The  most  remarkable  chemical  property 
of  hydrogen  peroxide  is  its  tendency  to  break  up  into  water  and 
oxygen,  as  represented  by  the  equation 

2H2O2=2H2O  +  O2, 

and  for  this  reason  it  is  a  fairly  energetic  oxidizing  agent.  The  rate 
at  which  the  peroxide  splits  up,  however,  depends  very  much  on  its 


i4o    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

purity  and  on  the  substances  with  which  it  is  brought  into  contact, 
As  already  mentioned,  a  solution  of  the  pure  peroxide  in  water  can 
be  concentrated  by  fractional  distillation  without  very  serious  decom- 
position taking  place,  but  an  impure  peroxide  cannot  be  concentrated 
in  this  way.  Many  finely-divided  substances,  such  as  silver,  gold, 
platinum  in  the  form  of  platinum  black,  manganese  dioxide,  and  even 
powdered  glass,  bring  about  a  very  rapid  decomposition  of  the  per- 
oxide. As  the  substances  remain  unchanged  at  the  end  of  the  re- 
action, we  have  to  do  with  catalytic  phenomena.  It  must  be  assumed 
that  under  all  circumstances  hydrogen  peroxide  is  slowly  decomposing 
and  that  the  rate  of  this  decomposition  is  greatly  increased  by  the 
catalysts  mentioned.  Even  the  rough  surfaces  in  ordinary  glass 
bottles  facilitate  the  decomposition  of  the  peroxide,  but  this  effect 
can  be  greatly  minimized  by  coating  the  interior  of  the  bottles  with 
solid  paraffin.  The  German  firm  Merck  supply  a  perfectly  pure  30 
per  cent,  aqueous  solution  of  hydrogen  peroxide  in  paraffin-coated 
bottles,  which  keeps  its  strength  remarkably  well.  The  peroxide  is 
very  unstable  in  alkaline  solution,  but  traces  of  acid  retard  decom- 
position very  markedly.  The  rate  of  decomposition  is  increased  by 
raising  the  temperature  and  also  by  exposure  to  light. 

In  the  cases  just  mentioned  free  oxygen  is  liberated,  but  many 
substances  take  up  an  atom  of  oxygen  from  the  peroxide  and  become 
oxidized,  whilst  water  remains.  Thus  black  lead  sulphide,  PbS,  is 
changed  by  hydrogen  peroxide  to  lead  sulphate,  PbSO4,  which  is 
white.  From  an  aqueous  solution  of  potassium  iodide  free  iodine 
is  liberated  and  at  once  turns  starch  emulsion  blue  (compare  ozone, 
p.  136),  the  equation  being  as  follows  :  — 


When  hydrogen  peroxide  is  added  to  a  solution  of  barium  or  calcium 
hydroxide  the  corresponding  peroxide  is  formed,  thus  : 

Ba(OH)2  +  H2O2->BaO2 


A  comparison  of  the  respective  formulae  shows  that  barium  peroxide 
may  be  regarded  as  being  formed  from  hydrogen  peroxide  by  putting 
in  an  atom  of  barium  for  two  atoms  of  hydrogen,  just  as  barium 
chloride,  BaCl2,  is  derived  from  two  molecules  of  hydrochloric  acid. 
From  this  point  of  view  barium  peroxide  may  be  regarded  as  being 
a  salt  derived  from  hydrogen  peroxide.  Of  the  different  compounds 
corresponding  with  the  general  formulae  MO2  and  M2O2  (M  =  metal), 
only  those  are  termed  true  peroxides  which  yield  hydrogen  peroxide 
on  treatment  with  acids  (cf.  p.  274). 


OZONE   AND   HYDROGEN    PEROXIDE          141 

As  hydrogen  peroxide  is  an  oxidizing  agent,  it  can,  of  course,  be 
used  for  bleaching  purposes.  It  can  be  safely  employed  for  bleaching 
very  delicate  materials,  such  as  silk,  hair,  and  ivory,  as  the  only 
products  are  water  and  oxygen.  For  the  same  reasons  it  is  a  valuable 
antiseptic. 

It  is  a  remarkable  fact  that  hydrogen  peroxide  sometimes  acts  as  a 
reducing  agent,  that  is,  it  removes  oxygen  from  other  substances.  It 
does  not  itself  become  more  highly  oxidized,  however,  but  gives  up 
an  atom  of  oxygen  which  with  a  further  atom  from  the  other  sub- 
stance forms  a  molecule  of  oxygen,  which  is  liberated.  Thus  it 
reduces  silver  oxide,  Ag2O,  to  metallic  silver  according  to  the  equation 


and  it  reduces  an  acidified  solution  of  potassium  permanganate, 
KMnO4,  to  colourless  salts.  The  latter  reaction  can  readily  be 
performed  by  adding  hydrogen  peroxide  to  a  colution  of  potassium 
permanganate  acidified  with  sulphuric  acid  ;  a  vigorous  evolution  of 
oxygen  will  be  noticed  and  the  original  deep  purple  colour  of  the 
solution  will  be  discharged. 

The  equation  expressing  this  change  is  rather  complicated  for  the 
present  stage  of  our  work,  but  will  be  better  understood  if  taken  in 
two  stages. 

(i) 
(2) 

We  may  suppose  that  potassium  permanganate  and  sulphuric  acid 
tend  to  react  according  to  the  first  equation,  giving  rise  to  potassium 
sulphate,  K2SO4,  manganous  sulphate,  MnSO4  (p.  546),  water  and 
oxygen,  and  that  this  tendency  becomes  effective  in  the  presence 
of  hydrogen  peroxide,  which  is  able  to  remove  the  oxygen,  as  repre- 
sented in  (2). 

When  ozone  and  hydrogen  peroxide  are  brought  together  in  not 
too  dilute  solution,  water  and  oxygen  are  formed  according  to  the 
equation 


Similarly  with  sodium  hypochlorite,  sodium  chloride  and  oxygen 
are  obtained  : 

NaOCl  +  H202->NaCl  +  H2O  +  O,. 


142     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

These  reactions  are  usually  explained  as  being  due  to  the  tend 
ency  of  the  loosely-held  atoms  in  the  peroxide  and  the  other  sub' 
stance  to  combine  with  formation  of  molecular  oxygen,  as  indicated 
above. 

Estimation  of  Hydrogen  Peroxide  Solutions — The  reaction  between  potas- 
sium permanganate  and  hydrogen  peroxide  in  acid  solution,  just  referred  to,  may 
be  applied  to  the  estimation  of  solutions  of  hydrogen  peroxide,  the  liberated 
oxygen  being  collected  and  measured  over  water  or  mercury.  The  apparatus 
used  for  this  purpose  is  the  same  as  that  already  described  in  connexion  with 
equivalents  (Fig.  40).  A  measured  (or  weighed)  quantity  of  the  solution  con- 
tained in  the  small  tube  is  placed  as  shown  in  the  bottle  B,  which  contains  excess 
of  potassium  permanganate  and  dilute  sulphuric  acid  and  is  connected  as  shown 
with  the  graduated  measuring  tube,  which  contains  water  (or  mercury).  The 
position  of  the  liquid  in  the  measuring  tube  is  adjusted  to  zero  on  the  scale 
at  atmospheric  pressure  by  means  of  the  levelling  tube,  the  pinchcock  being 
momentarily  opened  for  this  purpose.  The  levelling  tube  is  then  lowered,  and 
the  apparatus  tilted  so  that  the  contents  of  the  small  tube  become  thoroughly 
mixed  with  the  permanganate  solution.  When  the  reaction  is  over,  the  liquid  is 
brought  to  the  same  level  in  both  tubes,  and  the  volume  of  gas  read  off.  As  the 
original  air  in  the  apparatus  is  at  atmospheric  pressure  both  before  and  after 
the  experiment,  the  extra  volume  is  that  of  the  oxygen.  From  these  data  the  volume 
of  the  oxygen  at  o°,  and  hence  the  weight  of  oxygen  (or  of  hydrogen  peroxide) 
in  the  quantity  of  solution  taken,  may  readily  be  calculated.1 

The  concentration  of  a  solution  of  hydrogen  peroxide  is  often  expressed  in 
terms  cf  the  volume  of  oxygen  given  off  under  these  conditions  ;  thus  a  ao-volume 
solution  is  one  of  which  i  c.c.  yields  20  c.c.  of  oxygen  at  N.T.P.  Since  i  mol 
(34  grams)  of  hydrogen  peroxide  gives  with  permanganate  32  grams =22,4000.0. 

of  oxygen  at  N.T.P. ,  a  2o-volume  solution  contains  — — —  x  34=0.03  grams  in  i 

c.c.  of  solution,  i.e.  the  solution  contains  a  little  over  3  per  cent,  of  hydrogen 
peroxide.  < 

If  the  same  quantity  of  hydrogen  peroxide  were  decomposed  by  heat  alone, 
it  would  only  yield  half  the  volume  of  oxygen  obtained  by  the  use  of  perman- 
ganate. It  follows  that  a  ao-volume  solution  expressed  on  the  former  basis 
would  be  double  the  strength  of  a  2o-volume  solution  measured  by  means  of 
permanganate. 

Tests — (i)  A  delicate  test  for  hydrogen  peroxide  depends  upon 
the  fact  that  when  a  little  of  it  is  added  to  an  acidified  dilute  solution 
of  potassium  dichromate  a  deep  azure-blue  solution  is  obtained  (see 
p.  539).  A  dilute  solution  of  the  peroxide  is  shaken  up  with  ether,  a 
small  quantity  o£an  acidified  solution  of  potassium  dichromate  added, 

1  If  considerable  accuracy  is  required,  account  must  be  taken  of  the  fact  that 
the  oxygen  is  saturated  with  aqueous  vapour.  The  actual  pressure  of  the  oxygen 
is  that  of  the  atmosphere  less  the  pressure  of  aqueous  vapour  at  the  temperature 
of  the  experiment. 


THERMOCHEMISTRY  143 

and  after  further  shaking  the  upper  layer  of  ether,  which  separates 
on  standing  for  a  short  time,  is  coloured  deep  blue.  If  no  ether  is 
used,  the  blue  compound  rapidly  decomposes  in  aqueous  solution. 

(2)  The  most  delicate  test  for  hydrogen  peroxide  depends  upon  the 
production  of  a  coloured  solution  (orange-red  in  moderately  concen- 
trated solution,  yellow  in  very  dilute  solution)  when  the  peroxide  is 
added  to  a  colourless  solution  of  titanium  dioxide,  TiO2,  in  sulphuric 
acid.     The  coloured  substance  is  titanium  trioxide,  TiO3.     One  part 
of  hydrogen  peroxide  in  ten  million  parts  of  water  can  be  detected 
by  this  test. 

(3)  The  tests  just  described   are  characteristic  for  the  peroxide. 
Other  reactions,  such  as  the  liberation  of  iodine  from  an  acidified 
solution  of  potassium  iodide,  already  referred  to,  are  also  brought 
about  by  other  oxidizing  agents,  such  as  ozone.     Hydrogen  peroxide 
does  not  affect  an  emulsion  of  potassium  iodide  and  starch  in  the 
absence  of  acid,  but   at   once   turns   it  blue   when   a  little  ferrous 
sulphate  is  added. 

Graphic   Formulae  of  Ozone   and  of   Hydrogen   Per- 
oxide —  The  graphic  formula  of  ozone  may  be  written  thus  — 


O  -  O 

in  which  all  the  oxygen  atoms  are  divalent,  but  the  alternative 
formula  O  =  O  =  O,  in  which  one  of  the  oxygen  atoms  is  quadri- 
valent, appears  to  represent  better  the  readiness  with  which  an  atom 
of  oxygen  is  split  off,  leaving  a  molecule  of  oxygen  O  =  O.  That 
oxygen  can  act  under  certain  conditions  as  a  tetrad  has  been  estab- 
lished by  Collie  and  others  from  the  behaviour  of  certain  organic 
compounds. 

The  formula  for  hydrogen  peroxide  may  be  represented  thus  — 
H-O-OH,  in   which   the  oxygen   atoms  are  divalent,  or  thus- 

TJ 

pj>O  =  O,  one  of  the  oxygens  being  divalent  and  the  other  quadri- 

valent. The  latter  formula  is  perhaps  more  in  harmony  with  the 
great  tendency  of  the  peroxide  to  split  up  into  water  and  oxygen. 
Briihl,  on  the  basis  of  optical  measurements,  has  suggested  the 
formula  H-O  =  O-H,  both  oxygens  acting  as  tetrads. 

THERMOCHEMISTRY 

General  —  It  has  been  pointed  out  at  a  very  early  stage  of  our 
work  that  chemical  changes  are  invariably  associated  with  energy 


144     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

changes  in  the  system.  It  very  often  happens  that  heat  and  other 
forms  of  energy  are  given  out  as  the  result  of  chemical  changes, 
and  examples  have  been  met  with  in  the  combination  of  hydrogen 
and  oxygen  to  form  water,  and  in  the  combination  of  hydrogen  and 
chlorine  to  form  hydrogen  chloride.  On  the  basis  of  the  law  of  the 
conservation  of  energy,  we  state  in  such  cases  that  the  free  elements 
possess  more  chemical  energy  than  the  resulting  compounds,  and 
that  the  chemical  changes  are  accompanied  by  a  simultaneous  trans- 
formation of  chemical  energy  to  an  equivalent  quantity  of  other 
forms  of  energy,  mainly  heat.  The  department  of  chemistry  which 
is  concerned  with  the  heat  equivalent  of  chemical  changes  is  called 
thermochemistry. 

In  representing  the  results  of  thermochemical  measurements,  it  is 
convenient  to  deal,  not  with  equal  weights  of  substances,  but  with 
quantities  which  are  chemically  comparable,  that  is,  with  molecular 
or  molar  amounts.  Further,  when  heat  is  given  out  in  chemical 
change,  this  is  conveniently  represented  by  writing  the  number  of 
calories,  preceded  by  a  +  sign,  after  the  equation  representing  the 
chemical  change.  Thus  the  equation 

2H2O2->2H2O-|-O2  +  2  x  23,000  cal. 

indicates  that  when  68  grams  (twice  the  molar  weight)  of  hydrogen 
peroxide  split  up  into  water  and  oxygen  gas,  2  x  23,000  calories 
are  given  out ;  in  other  words,  the  chemical  energy  associated 
with  68  grams  of  hydrogen  peroxide  is  greater  by  46,000  cal.  than 
that  associated  with  its  products  of  decomposition.  When  heat  is 
absorbed  in  a  chemical  change,  the  amount  so  absorbed  is  preceded 
by  the  -  sign.  Thus  the  equation 

2  HC1->H2+C12- 44,000  cal. 

indicates  that  in  splitting  up  73  grams  of  hydrochloric  acid  into 
2  grams  of  hydrogen  and  71  grams  of  chlorine,  44,000  calories  are 
taken  up  ;  in  other  words,  the  chemical  energy  associated  with 
73  grams  of  hydrochloric  acid  is  44,000  calories  less  than  that 
associated  with  its  products  of  decomposition. 

When  a  definite  amount  of  heat  is  given  out  in  a  chemical  change, 
according  to  the  law  of  the  conservation  of  energy,  exactly  the  same 
amount  must  be  supplied  in  order  to  regain  the  original  substances 
in  the  same  amounts.  Thus  the  thermochemical  equation  represent- 


THERMOCHEMISTRY  145 

ing  the  combination  of  hydrogen  and  chlorine  to  hydrochloric  acid 
is  as  follows  : — 

H2+C12->2HC1 +  44,000  cal. 

As  the  heat  is  given  out  when  the  change  proceeds  in  the  direction 
of  the  arrow,  the  +  sign  is  used. 

The  heat  given  out  or  absorbed  when  a  mol  of  substance  is  formed 
from  its  elements  is  termed  its  heat  of  formation  ;  the  heat  given  out 
or  absorbed  when  a  mol  of  substance  is  split  up  into  its  component 
elements  is  termed  its  heat  of  decomposition.  Most  substances,  like 
hydrogen  chloride  and  water,  are  formed  with  evolution  of  heat,  and 
their  heat  of  formation  is  said  to  be  positive,  a  few,  such  as  ozone 
(p.  135)  and  the  oxides  of  chlorine  (p.  177)  are  formed  with  absorption 
of  heat,  so  that  their  heat  of  formation  is  negative.  The  thermo- 
chemical  equation  for  the  formation  of  ozone  is  as  follows  : — 

3O2->2O8  — 2  x  34,100  cal. 

Chemical  compounds  formed  from  their  elements  with  evolution  of 
heat  are  said  to  be  exothermic,  while  substances  formed  from  their 
elements  with  absorption  of  heat  are  endothermic.  It  is,  of  course, 
evident  that  heat  must  be  supplied  to  split  up  exothermic  substances, 
whereas  endothermic  substances  decompose  with  evolution  of  heat. 

A  little  consideration  will  show  that,  when  the  other  conditions 
are  kept  constant,  rise  of  temperature  favours  the  decomposition 
of  exothermic  compounds,  but  favours  the  formation  of  endothermic 
compounds  from  their  components  (Le  Chatelier's  theorem,  p.  171). 

Two  other  terms  employed  in  thermochemical  work  may  be  men- 
tioned. The  heat  of  combustion  of  a  substance  is  that  amount  of  heat 
given  out  when  a  mol  of  it  is  completely  burned,  and  the  term  is 
generally  applied  to  combustion  in  oxygen.  Heat  of  solution  is 
the  heat  given  out  when  a  mol  of  a  substance  is  dissolved  in  a  large 
excess  of  the  solvent. 

Energy  Content  of  Different  Forms  of  Substances — 
It  is  very  important  in  writing  thermochemical  equations  to  see  that 
the  particular  forms  of  the  substances  taking  part  in  the  reactions 
are  clearly  stated,  as  these  may  differ  considerably  in  their  energy 
content.  When  hydrogen  and  oxygen  unite  to  form  liquid  water, 
68,400  cal.  per  mol  are  given  out.  Part  of  this  heat,  however,  is 
due  to  the  change  of  gaseous  to  liquid  water,  and  as  the  heat  of 
vaporization  is  537  cal.  per  gram  at  100°  (p.  66),  the  heat  per  mol 
due  to  the  change  of  state  is  537  x  18  =  9666  calories.  It  follows  that 

10 


146     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  heat  given  out  when  hydrogen  and  oxygen  unite  to  form  a  mol 
of  steam  at  100°  is  about  68,400-9670  =  58,730  cal.  The  difference 
in  the  energy  content  of  water  and  ice  at  o°  is,  on  the  same  basis, 
Sox  18=1440  cal.  per  mol. 

A  further  important  correction  has  to  be  taken  into  consideration 
when  gases  are  formed  or  disappear  as  the  result  of  a  chemical 
change.  When  a  gas  is  generated  under  atmospheric  pressure,  it 
does  work  and  heat  is  taken  up  ;  when,  on  the  other  hand,  a  gas 
disappears  heat  is  given  out.  When  a  mol  of  any  gas  is  formed 
against  external  pressure  at  the  absolute  temperature  T,  the  heat 
absorbed  is  2T,  and  when  a  mol  of  any  gas  disappears  2T  calories 
are  given  out.  It  must  be  carefully  noted  that  these  heat  changes 
are  in  addition  to  those  associated  with  the  chemical  changes.  It 
follows  from  the  above  that  in  the  reaction 

2H2  +  Q2  =  2H2O  (liquid)  +2x68,400  cal., 

the  observed  heat  of  combustion  includes  3X2T  calories  =1640  cal. 
(taking  the  temperature  of  the  experiment  as  o°  =  273  abs.)  due  to 
the  disappearance  of  three  mols  of  gas.  When  the  change  is  carried 
out  under  such  conditions  that  the  volume  is  kept  constant  no  external 
work  can  be  done.  The  above  considerations  enable  us  to  calculate 
the  difference  in  the  thermal  effect  under  constant  volume  and  constant 
pressure  conditions  respectively.  \ 

It  is  evident  from  the  case  of  ozone  and  oxygen  that  the  different 
allotropic  modifications  of  a  substance  may  differ  considerably  in 
energy  content,  and  many  illustrations  of  this  will  be  met  with  later. 

Hess's  Law — A  very  important  law  of  thermochemistry,  first 
established  experimentally  by  Hess,  may  be  expressed  as  follows : 
When  a  chemical  change  takes  place  between  definite  amounts  of 
different  substances,  the  amount  of  heat  given  out  is  always  the 
same  provided  the  initial  and  final  products  are  the  same  in  each 
case.  As  an  illustration  of  this  law  we  will  consider  the  combination 
of  carbon  and  oxygen  to  form  carbon  dioxide.  This  reaction  may 
take  place  slowly  or  quickly,  and  may  take  place  in  one  or  more 
stages,  but  in  all  cases,  if  one  starts  with  12  grams  of  carbon  and  32 
grams  of  oxygen,  and  finishes  up  with  carbon  dioxide  alone,  94,300 
cal.  are  always  liberated.  The  temperature  attained  may,  of  course, 
be  very  different  according  to  the  speed  of  combination.  When  the 
combustion  is  slow,  the  heat  evolved  is  communicated  to  the  sur- 
roundings and  conducted  away,  and  the  temperature  attained  is  much 
lower  than  when  the  combustion  is  rapid.  The  total  amount  of  heat 
given  out  is,  however,  the  same  in  both  cases. 


THERMOCHEMISTRY  147 

The  chief  importance  of  Hess's  law  is  that  it  enables  us  to  deter- 
mine the  thermal  equivalent  of  many  reactions  which  cannot  readily 
be  carried  out  directly.  Suppose,  for  example,  we  wish  to  determine 
the  heat  given  out  or  absorbed  when  12  grams  of  carbon  combines 
with  4  grams  of  hydrogen  to  form  marsh  gas,  CH4  (p.  345). 

We  can  assume  that  the  reaction  is  carried  out  in  one  stage 

CH4  +  2O2=CO2  +  2H2O  +  2i  1,900  cal.  (a) 

or  in  the  following  two  stages 

cal.  (b) 


l.      1 
cal.  J 


As,  by  Hess's  law,  the  heat  given  out  in  the  reaction  (a)  is  neces- 
sarily the  same  as  that  given  out  when  the  same  reaction  takes  place 
in  two  stages  (b}  and  (c),  we  have 

2  1  1  ,900  =  x  +  94,300  +  1  36,800, 

when  x=  -  19,200  cal. 

As  19,200  cal.  are  absorbed  when  methane  is  split  up  into  its 
elements,  the  same  amount  is  given  out  when  it  is  formed  from  its 
elements  ;  in  other  words,  methane  is  an  exothermic  compound. 

Relationship  between  Chemical  Reactivity  and  Heat 
of  Reaction  —  It  will  be  evident,  on  consideration  of  the  reactions 
already  discussed,  that  chemical  changes  may  be  divided  into  two 
classes:  (i)  those  which  under  the  conditions  of  experiment  are 
spontaneous  or  proceed  of  themselves  once  they  are  started,  (2) 
those  which  only  proceed  when  forced  by  some  external  energy,  for 
example,  when  heat  is  continuously  supplied.  As  examples  of  the 
first  class  we  have  the  combination  of  hydrogen  and  oxygen  to  form 
water  and  of  hydrogen  and  chlorine  to  form  hydrogen  chloride  ;  as 
examples  of  the  second  class  the  splitting  up  of  mercuric  oxide  and 
of  potassium  chlorate  by  heat.  The  former  reactions  proceed  rapidly 
to  completion  with  explosive  violence,  the  latter  stop  at  once  when 
heating  is  stopped. 

The  energy  relations  of  the  systems  throw  a  great  deal  of  light  on 
these  changes.  We  have  already  seen  that  the  formation  of  water 
and  of  hydrogen  chloride  from  the  respective  elements  are  highly 
exothermic  reactions  (p.  145).  On  the  other  hand,  the  thermochemicai 
equation  for  the  decomposition  of  mercuric  oxide  is  as  follows  :  — 
2HgO-»2Hg  +  O2  -  2  x  20,600  cal. 


148     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

So  that  a  large  amount  of  heat  has  to  be  supplied.  From  these 
results  the  conclusion  might  be  drawn  that  a  chemical  reaction  pro- 
ceeds of  itself  in  the  direction  in  which  heat  is  given  out,  that  ts,  that 
spontaneous  chemical  changes  are  exothermic.  Experience  has  shown 
that  this  is  actually  the  case  for  a  very  large  number  of  chemical 
changes,  more  particularly  those  with  large  heats  of  reaction,  and 
the  above  statement  may  be  used  as  an  approximate  rule  by  the 
student.  It  does  not  apply  in  every  case,  however,  as  there  are 
many  changes,  for  example,  the  dissolution  of  ammonium  chloride 
and  certain  other  salts  in  water,  which  proceed  of  themselves  and 
produce  a  considerable  cooling  effect.  It  would  lead  too  far  to  discuss 
these  relationships  fully,1  and  it  is  quite  sufficient  for  general  purposes 
to  make  use  of  the  approximate  rule  already  mentioned. 

A  reaction  which  is  endothermic  or  only  feebly  exothermic,  and 
which  scarcely  proceeds  at  all  under  ordinary  conditions,  may  often 
be  brought  about  by  associating  it  with  some  other  change,  so  that 
the  total  change  is  now  strongly  exothermic.  The  reaction  usually 
adduced  in  illustration  of  this  statement  is  that  between  hydrogen 
sulphide  and  iodine  (p.  161).  When  dry  hydrogen  sulphide  is  passed 
over  dry  iodine  at  room  temperature  no  appreciable  reaction  occurs. 
As  a  matter  of  fact,  the  reaction  is  endothermic : 

H2S  + 12->2  H I  +  S  -  7300  cal. 

When,  however,  hydrogen  sulphide  is  passed  into  iodine  suspended 
in  excess  of  water  at  room  temperature  the  reaction  occurs  readily, 
hydriodic  acid  and  sulphur  being  formed.  Taking  into  account  the 
fact  that  the  hydrogen  iodide  is  finally  present  in  aqueous  solution  * 
the  thermochemical  equation  now  becomes 

H2S,Aq  +  I2-»2HI,Aq  +  S  + 17,000  cal., 

that  is,  the  total  change  is  now  strongly  exothermic  and  the  reaction 
proceeds  spontaneously.  The  alteration  in  the  thermal  character 
of  the  reaction  is  due  in  this  case  to  the  high  heat  of  solution  of 
hydriodic  acid  (p.  163). 

Other  illustrations  of  this  important  principle  will  be  met  with  in 
the  course  of  our  work. 

1  Cf.  Physical  Chemistry,  p.  148. 

2  When  a  substance  in  dilute  aqueous  solution  takes  part  in  a  chemical  change, 
this  is  sometimes  indicated  by  appending  Aq  to  the  formula  of  the  substance,  as 
above. 


CHAPTER  XII 
THE   HALOGENS  AND   HALOGEN   ACIDS 

IN  Chapter  VIII.  chlorine  and  its  compound  with  hydrogen, 
hydrogen  chloride,  have  been  considered.  In  the  present  chapter 
we  shall  deal  with  three  elements — fluorine,  bromine,  and  iodine — 
which  have  so  many  analogies  with  chlorine,  both  as  regards  the 
elements  themselves  and  their  more  important  compounds,  that  all 
four  elements  are  said  to  belong  to  the  same  family.  They  are  called 
halogens  (from  a\s  salt,  and  ycwato,  I  produce)  in  allusion  to  common 
salt,  the  chief  source  of  the  chlorine  compounds. 

Like  chlorine,  the  three  other  elements  each  form  a  compound  with 
hydrogen  of  the  same  type  as  hydrogen  chloride ;  the  formulae  and 
names  are  as  follows :  Hydrogen  fluoride,  HF ;  hydrogen  bromide, 
HBr  ;  hydrogen  iodide,  HI.  All  these  compounds  are  gases  at  the 
ordinary  temperature  (hydrogen  fluoride  boils  at  19.5°),  and  dissolve 
in  water  to  form  strong  acids.  These  hydrogen  compounds  are  also 
described  in  this  chapter. 

FLUORINE 

Symbol,  F.     Atomic  weight,  19.     Molecular  weight,  38. 

Occurrence— The  chief  source  of  fluorine  compounds  is  calcium 
fluoride,  CaF2,  which  occurs  naturally  as  fluor-spar  or  Derbyshire 
spar.  The  pure  mineral  forms  colourless  cubical  crystals,  but  many 
samples  are  brilliantly  coloured  by  traces  of  impurities.  Another 
naturally  occurring  compound  of  fluorine  is  cryolite,  AlF3,3NaF,  a 
double  fluoride  of  sodium  and  aluminium.  Fluorine  occurs  to  a  small 
extent  in  bones  and  in  the  enamel  of  teeth. 

Preparation — It  is  an  interesting  fact  that  fluorine  was  not 
obtained  in  the  free  condition  till  1886.  We  have  learnt  that  when 
hydrogen  chloride  is  electrolyzed,  part  of  the  chlorine  acts  on  the 
water  and  sets  free  oxygen.  Fluorine  has  a  still  greater  affinity  for 
hydrogen  than  chlorine  has,  and  when  an  aqueous  solution  of  hydrogen 
fluoride  is  electrolyzed  the  fluorine  set  free  at  the  anode  acts 

149 


i5o     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

immediately  on  the  water,  with  liberation  of  oxygen  and  reformation 
of  hydrogen  fluoride.  It  appears  at  first  sight  as  if  this  difficulty  could 
be  overcome  by  using  anhydrous  hydrogen  fluoride,  but  Gore  found  that 
the  latter  does  not  conduct  the  electric  current,  so  that  no  chemical 
change  occurs.  Moreover,  fluorine  acts  energetically  on  the  great 
majority  of  elements  and  of  chemical  compounds,  so  that  it  is  not 


FIG.  43. 

surprising  that  the  attempts  made  to  isolate  this  element  were  at 
first  unsuccessful. 

The  problem  was  solved  by  Moissan,  who  subjected  to  electrolysis 
anhydrous  hydrogen  fluoride  in  which  potassium  fluoride  had  been 
dissolved  to  make  a  conducting  solution.  One  form  of  apparatus 
used  is  represented  in  Fig.  43.  The  U-tube  (made  of  an  alloy  of 
platinum  and  iridium),  which  contains  the  mixture  of  hydrogen 
fluoride  and  potassium  fluoride  (about  4  parts  to  i  by  weight)  is 
provided  with  two  side  tubes,  A  and  B,  and  the  open  ends  are  closed 
with  stoppers,  C,  made  of  fluor-spar  and  wrapped  in  thin  sheet 


THE   HALOGENS   AND    HALOGEN   ACIDS        151 

platinum.  The  electrodes,  which  are  also  made  of  a  mixture  of  platinum 
and  iridium  (one  of  the  few  materials  which  are  not  much  affected 
by  fluorine),  pass  through  the  stoppers,  and  are  kept  in  place  by 
screws. 

During  the  electrolysis  the  apparatus  is  kept  at  a  temperature  of 
-  23°  by  means  of  boiling  methyl  chloride.  The  fluorine,  which  is 
given  off  as  a  gas  at  the  positive  pole,  is  passed  through  a  platinum 
spiral  kept  at  A  low  temperature  in  order  to  remove  any  hydrogen 
fluoride,  and  may  then  be  collected  and  examined  in  a  platinum  tube, 
the  ends  of  which  are  closed  with  plates  of  fluor-spar. 

In  his  earlier  experiments  Moissan  used  an  apparatus  of  platinum- 
iridium,  but  found  subsequently  that  an  apparatus  of  copper  was 
equally  convenient  and  much  less  costly. 

The  chemical  changes  may  be  assumed  to  take  place  mainly  in 
accordance  with  the  equations 

(i) 
(2) 


The  potassium  liberated  at  the  cathode  as  the  result  of  the  primary 
change  represented  by  equation  (i)  immediately  reacts  with  hydrogen 
fluoride,  liberating  hydrogen  and  reforming  potassium  fluoride. 

Physical  Properties  —  At  ordinary  temperatures  fluorine  is 
a  greenish-yellow  gas,  much  lighter  in  colour  than  chlorine  ;  it  has 
a  very  pungent  odour.  The  density  of  the  gas  is  approximately  19, 
corresponding  with  the  molecular  formula  F2.  Fluorine  has  been 
obtained  by  Moissan  and  Dewar  as  a  bright  yellow  liquid  boiling  at 
—  187°  ;  the  liquid  is  almost  without  action  on  glass.  On  cooling 
with  liquid  hydrogen,  fluorine  forms  a  pale  yellow  solid  melting  at 
-223°,  and  on  further  cooling  to  —252°  the  solid  becomes  perfectly 
white. 

Chemical  Properties  —  Fluorine  is  the  most  chemically  active 
element  known  ;  it  combines  directly  with  all  other  elements  except 
oxygen  and  the  elements  of  the  helium  family,  and  in  many  cases 
with  explosive  violence.  It  combines  with  hydrogen  explosively  in 
the  dark  at  ordinary  temperatures,  and  even  at  —210°  the  gases 
unite  immediately,  producing  a  flame.  Sulphur,  phosphorus  and 
iodine  melt  and  then  inflame  in  fluorine  ;  arsenic,  antimony  and 
boron  become  incandescent,  and  crystalline  silicon  burns  in  it  with 
great  brilliancy.  Fluorine  does  not  combine  with  oxygen  under  any 
conditions  so  far  realized. 

Most  of  the  metals  combine  readily  with  fluorine  at  the  ordinary 


152     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

temperature,  but  gold  and  platinum  are  only  slightly  affected,  much 
less  than  when  acted  on  by  chlorine.  With  copper  a  thin  coating  of 
the  fluoride  is  formed,  which  protects  the  metal  against  further  action. 
Glass,  whether  moist  or  dry,  is  at  once  attacked  by  fluorine. 

Owing  to   its  great   affinity  for  hydrogen,  fluorine  immediately 
decomposes  water  according  to  the  equation 


The  liberated  oxygen  contains  a  large  proportion  (up  to  12  per  cent.) 
of  ozone. 

HYDROFLUORIC  ACID  (HYDROGEN  FLUORIDE),  HF 

Preparation  —  (i)  Hydrogen  fluoride  is  most  conveniently  ob- 
tained by  heating  calcium  fluoride  (fluor-spar),  CaF2,  with  sulphuric 
acid  in  a  vessel  of  lead  or  platinum,  as  glass  is  readily  acted  on  by 
the  fluoride.  The  equation  is  as  follows  :  — 

CaF2+H2SO4-^CaSO4 


The  volatile  fluoride  is  absorbed  in  water,  and  the  aqueous  solution, 
which  is  termed  hydrofluoric  acid,  is  kept  in  lead,  paraffin,  or  rubber 
vessels. 

(2)  Anhydrous  hydrogen  fluoride  is  best  prepared  by  heating  dry 
hydrogen  potassium  fluoride,  HF,KF,  in  an  apparatus  of  platinum, 
the  condenser  and  receiver  being  surrounded  by  ice.  The  double 
compound  is  decomposed  on  heating,  according  to  the  equation 

HF,KF->KF  +  HF, 

and  the  hydrogen  fluoride  is  collected  in  the  cooled  receiver  as 
a  colourless  liquid. 

Hydrogen  fluoride  may  also  be  prepared  by  other  methods 
analogous  to  those  described  under  hydrogen  chloride  (q.v.}. 

Physical  Properties  —  Hydrogen  fluoride  is  a  colourless, 
fuming  liquid,  boiling  at  19.5°.  The  vapour  is  very  pungent,  and 
when  inhaled  has  a  very  injurious  effect  on  the  mucous  membrane. 
It  is  usually  kept  in  platinum  vessels,  but  when  perfectly  free  from 
water  it  has  no  action  on  glass  (Gore).  It  mixes  with  water  in  all 
proportions,  and  when  the  aqueous  solution  Is  distilled  a  mixture  of 
constant  boiling-point  (120°  at  760  mm.  pressure)  containing  36  per 
cent,  of  hydrogen  fluoride  is  obtained. 

At  88°  the  vapour  density  of  hydrogen  fluoride  is  approximately  10, 


THE   HALOGENS   AND   HALOGEN   ACIDS        153 

corresponding  with  the  formula  HF  (molecular  weight  =  20),  but  as 
the  temperature  is  lowered  the  density  progressively  increases,  and 
at  26.4°  is  25.6,  corresponding  with  a  molecular  weight  of  51.2.  The 
simplest  explanation  of  this  remarkable  fact  would  appear  to  be  that 
as  the  temperature  falls  the  simple  HF  molecules  unite  among  them- 
selves to  form  compounds  of  the  type  (HF)^,  where  x  is  a  whole 
number.  If  the  whole  of  the  fluoride  was  present  as  H2F2  molecules, 
the  molecular  weight  would  be  40,  but  this  number  is  already 
exceeded  at  26°,  so  that  some  of  the  molecules  must  be  still  more 
complex.  In  such  a  case  the  substance  is  said  to  associate  or  poly- 
merize as  the  temperature  falls.  Many  other  cases  of  association  are 
known,  and  they  will  be  fully  considered  in  the  next  chapter  under 
the  heading  of  Chemical  Equilibrium. 

Chemical  Properties—  As  in  the  case  of  hydrogen  chloride,  a 
clear  distinction  must  be  drawn  between  hydrogen  fluoride,  which 
has  no  acid  properties,  and  a  mixture  of  the  fluoride  with  water, 
which  is  a  typical  acid,  hydrofluoric  acid.  The  latter  acid  attacks 
many  metals,  liberating  hydrogen  and  forming  fluorides.  The  for- 
mulae of  the  fluorides  correspond  with  those  of  the  chlorides.  Most 
of  them  are  soluble  in  water,  but  the  fluorides  of  calcium  and  barium 
are  insoluble. 

The  most  striking  property  of  hydrofluoric  acid  is  that  it  acts  on 
glass,  and  it  is  therefore  largely  used  for  etching  purposes.  The 
object  to  be  etched  is  covered  with  wax,  and  the  figures  or  other 
marks  cut  through  the  wax  with  a  pointed  instrument.  The  prepared 
surface  is  then  expose  "*  to  the  fumes  of  hydrogen  fluoride  for  a  time 
or  dipped  into  an  aqueous  solution  of  the  acid.  Only  the  parts  of  the 
glass  where  the  coating  is  removed  are  affected  by  the  acid,  and  on 
removing  the  rest  of  the  coating  the  design  will  be  found  etched  on 
the  glass. 

The  effect  just  described  depends  upon  the  action  of  the  hydro- 
fluoric acid  on  silicates  (of  which  glass  is  composed,  p.  454),  according 
to  the  equation 


the  silicon  fluoride,  SiF4,  which  is  the  main  product  of  the  action, 
escaping  as  a  gas  (p.  365). 

Acids  such  as  hydrochloric  acid  and  hydrofluoric  acid,  which  con- 
tain only  one  hydrogen  atom  replaceable  by  a  metal,  are  termed 
monobasic;  acids  containing  more  than  one  replaceable  hydrogen 
atom  are  termed  polybasic. 


154     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

BROMINE 

Atomic  weight,  79.92.     Molecular  weight,  159.84. 

History — Bromine  was  discovered  in  1826  by  Balard  in  the 
mother  liquor  obtained  by  removing  most  of  the  sodium  chloride  from 
sea-water.  The  name  is  derived  from  fipapos,  a  stench,  in  allusion  to 
its  irritating  odour. 

Occurrence — On  account  of  its  great  affinity  for  other  elements 
bromine  is  never  found  free  in  nature.  It  occurs  in  small  amount  in 
sea-water  in  combination  chiefly  with  sodium,  potassium,  and 
magnesium,  and,  owing  to  the  fact  that  bromides  are  very  soluble  in 
water,  it  remains  behind  in  the  mother  liquor  after  the  greater  part 
of  the  sodium  chloride  has  been  separated  by  crystallization.  It  also 
occurs,  mainly  as  magnesium  bromide,  associated  with  carnallite, 
MgCl2,KCl,6H2O,  a  double  chloride  of  magnesium  and  potassium,  in 
the  salt  deposits  at  Stassfurt,  and  this  is  now  the  chief  commercial 
source.  Bromides  are  also  met  with  in  certain  mineral  springs. 

Preparation — (i)  Bromine  is  now  usually  prepared  commercially 
by  the  action  of  chlorine  on  the  bromides  present  in  the  mother  liquors 
obtained  in  the   preparation  of  potassium   chloride   from  carnallite 
(p.  412),  or  in  the  preparation  of  common  salt  from  sea-water  : 
9  MgBr2+Cl2=MgCl2+Br2. 

This  method  of  preparation  depends  upon  the  fact  that  chlorine  is  a 
more  active  element  than  bromine,  and  displaces  the  latter  from 
combination. 

The  solution  containing  the  bromides  enters  the  top  of  a  tower 
lined  with  stoneware  and  provided  with  perforated  plates,  and  on  its 
way  down  meets  an  ascending  current  of  chlorine  which  liberates  the 
bromine.  The  latter,  in  the  form  of  vapour,  escapes  through  a  pipe 
near  the  top  of  the  tower  and  is  condensed  in  a  cooled  receiver.  In 
order  to  remove  traces  of  chlorine,  some  calcium  or  ferrous  bromide 
is  added  to  the  bromine  and  the  latter  then  redistilled. 

(2)  A  less  advantageous  method  of  obtaining  bromine  from  bromides 
(but  one  which  can  be  used  for  laboratory  purposes)  is  to  act  on  the 
latter  with  manganese  dioxide  and  sulphuric  or  hydrochloric  acid  : 

MgBr2+2H2SO4+MnO2=MnSO4+MgSO4+2H2 
MgBr2+4HCl  + MnO2-MnCl2+MgCl2+2H2 

On  the   commercial  scale  dilute  sulphuric  acid  is  added  to  the 


THE   HALOGENS   AND    HALOGEN   ACIDS       155 

mother  liquor  containing  bromides,  the  sulphates  which  separate  out 
on  standing  are  removed,  manganese  dioxide  and  sulphuric  acid  are 
added,  and  the  mixture  distilled.  The  impure  distillate  in  the  first 
receiver  is  rejected,  the  bromine  is  absorbed  in  sodium  hydroxide 
solution  in  the  second  receiver,  the  solution  evaporated,  ignited,  and 
again  distilled  with  manganese  dioxide  and  sulphuric  acid,  when  fairly 
pure  bromine  is  obtained.  A  still  purer  product  could  be  obtained  by 
repeated  recrystallization  of  the  alkali  bromide  before  redistillation. 

(3)  When  an  electric  current  is  passed  through  a  solution  of  a 
bromide,  bromine  is  set  free  at  the  anode.  This  method  is  now  used 
commercially  for  obtaining  bromine  directly  from  mixtures  of  bro- 
mides and  chlorides,  as  under  suitable  conditions  all  the  bromine  can 
be  liberated  before  the  chlorine  begins  to  separate. 

Physical  Properties—  Bromine  is  a  dark-red  liquid  at  ordinary 
temperatures.  Its  density  at  o°  is  3.188.  It  boils  at  59°,  and  solid 
bromine  melts  at  -7.3°.  It  has  a  high  vapour  pressure  at  room 
temperature,  giving  off  brown  fumes  which  have  a  very  irritating 
effect  on  the  nose  and  throat  and  also  on  the  eyes. 

Bromine  dissolves  fairly  readily  in  water,  forming  a  red  solution 
known  as  bromine  water;  100  grams  of  water  at  o°  dissolve  3.6 
grams,  and  at  20°  3.2  grams  of  bromine.  On  cooling  the  aqueous 
solution  a  crystalline  hydrate  of  bromine,  Br2,ioH2O,  is  obtained 
(cf.  chlorine,  p.  91). 

The  vapour  density  of  bromine  up  to  750°  is  about  80,  hence  the 
molecular  weight  is  about  160  ;  and  as  the  atomic  weight  of  bromine, 
the  smallest  quantity  of  it  which  occurs  in  a  molecule  referred  to  the 
atom  of  oxygen  as  16  (p.  116)  is  80,  the  molecular  formula  is  Br2.  It 
was  found  by  Victor  Meyer,  however,  that  at  higher  temperatures  the 
density  of  bromine  diminishes,  and  at  1500°  is  considerably  less  than 
80.  The  simplest  explanation  of  this  observation  is  that  at  high 
temperatures  bromine  is  partly  split  up  or  dissociated  with  single 
atoms,  according  to  the  equation 


This  important  point  will  be  further  referred  to  in  the  present  chapter 
(p.  159;.  _ 

Chemical  Properties  —  The  chemical  properties  of  bromine 
are  summarized  in  the  statement  that  it  behaves  like  chlorine,  but 
acts  less  energetically.  It  unites  directly  with  hydrogen,  but  the 
combination  does  not  proceed  explosively  in  sunlight,  as  in  the  case 
of  chlorine.  It  combines  directly  with  non-metals  like  phosphorus 


156     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

and  arsenic,  and  also  with  many  metals.  It  has  a  slight  bleaching 
action,  which  may  be  accounted  for  as  described  under  chlorine 
(p.  91).  The  great  chemical  activity  of  chlorine  as  compared  with 
bromine  is  further  shown  by  the  fact  already  mentioned,  that  the 
former  displaces  the  latter  from  combination. 


HYDROGEN  BROMIDE  (HYDROBROMIC  ACID),  HBr 

Preparation — (i)  Hydrogen  bromide  can  be  obtained  by  direct 
combination  of  its  elements.  Hydrogen  and  bromine  do  not  combine 
explosively  even  when  a  lighted  taper  is  applied  to  the  mixture,  but 
when  the  mixed  gases  are  passed  through  a  red-hot  tube  containing 
finely  divided  platinum  (catalytic  agent,  p.  22)  fumes  of  hydrogen 
bromide  are  produced.  Further,  hydrogen  burns  in  bromine  vapour, 
giving  rise  to  hydrogen  bromide. 

(2)  As  hydrogen  bromide  is  readily  volatile,1  it  can  be  obtained  by 
the  action  of  a  suitable  acid  ;  for  example,  phosphoric  acid,  H3PO4, 
on  a  bromide : 

KBr  +  HsPO4->KH2PO4-r-HBrt. 

(3)  It  might  appear  that  in  the  above  method  of  preparation  phos- 
phoric acid  could  advantageously  be  replaced  by  sulphuric  acid ;  but 
if  the  latter  acid  is  used,  it  will  be  observed  that  brown  fumes  of 
bromine  are  also  liberated.     The  reaction  in  this  case  takes  place  in 
two  stages : 

(1)  2KBr  +  H2S04->2HBr  +  K2S04, 

(2)  2HBr+H2SO4->2H2O  +  SO2  +  Br2, 

the  bromine  resulting  from  a  secondary  reaction,  in  which  the  sul- 
phuric acid  is  reduced  to  sulphur  dioxide  and  water,  and  the  hydrogen 
bromide  oxidized  to  bromine  and  water. 

(4)  The  most  satisfactory  method  for  the  preparation  of  hydrogen 
bromide  is  to  decompose  phosphorus  tribromide,  PBr3,  with  water: 

PBrs  +  sHOH->3HBr+  H3PO3. 

As  the  other  product  of  the  reaction,  phosphorous  acid,  H3PO3,  is 
non-volatile,  the  substances  can  readily  be  separated. 

In  practice,  the  phosphorus  bromide  is  formed  and  decomposed  in 

1  The  formation  of  a  volatile  product  is  conveniently  shown  by  an  arrow 
directed  upwards,  the  formation  of  a  precipitate  by  a  downwardly  directed 
arrow. 


THE    HALOGENS   AND    HALOGEN   ACIDS        157 

one  operation.  Red  phosphorus,  mixed  with  a  little  water,  is  placed 
in  the  flask  A  (Fig.  44),  the  latter  being  closed  by  a  cork  carrying  a 
dropping-funnel  containing  bromine.  The  bromine  is  added  drop  by 
drop  to  the  mixture,  and  the  issuing  gas  passed  through  a  U-tube  B 
containing  glass  beads  mixed  with  red  phosphorus,  in  order  to  remove 
any  free  bromine  which  may  be  carried  over.  The  hydrogen  bromide 
may  be  collected  in  a  gas  jar  by  upward  displacement  of  air  (p.  39), 
or  an  aqueous  solution  may  be  prepared  by  supporting  the  end  of  the 
delivery  tube  just  over  the  surface  of  water  in  a  bottle  C. 

(5)  An  aqueous  solution  of  hydrogen  bromide  is  readily  obtained 
by  passing  hydrogen  sulphide,  H2S,  through  bromine  water  till  the 
colour  of  the  latter  is  discharged,  and  then  removing  the  precipitated 
sulphur  by  filtration.  The  main  reaction  is  represented  by  the 
equation 


Physical  Properties  —  Hydrogen  bromide,  like  hydrogen 
chloride,  is  a  colourless  gas  with  a  sharp  odour,  and  fumes  in  con- 
tact with  moist  air.  It  can  be  condensed  to  a  liquid,  which  boils  at 
-68.7°,  and  on  further  cooling  solidifies  ;  the  crystals  melt  at  -80°. 
One  volume  of  water  dissolves  about  600  volumes  of  the  gas  at  10°. 
When  distilled,  hydrobromic  acid,  like  hydrochloric  acid,  forms  a 
mixture  of  constant  boiling-point  (126°  at  760  mm.  pressure)  contain- 
ing 48  per  cent,  of  hydrogen  bromide. 

The  thermochemical  equation  representing  the  heat  of  formation 
of  gaseous  hydrogen  bromide  from  its  components  in  the  gaseous 
form  is  as  follows  :  — 

H2+Br2=2HBr  +  2  x  12,100  cal. 

Further,  the  heat  of  solution  (p.  145)  of  hydrogen  bromide  is  20,000 
calories,  so  that  12,100  +  20,000=32,100  cal.  are  given  out  when 
I  gram  of  hydrogen  and  80  grams  of  bromine  combine  and  the 
product  is  dissolved  in  excess  of  water. 

Chemical  Properties  —  Liquefied  hydrogen  bromide  has  no 
acidic  properties,  but  the  solution  of  the  gas  in  water,  known  as 
hydrobromic  acid,  is  a  typical  acid.  It  acts  on  many  metals,  forming 
bromides  and  liberating  hydrogen.  Oxidizing  agents  set  free  bromine 
from  hydrobromic  acid,  a  change  which  takes  place  much  more 
readily  than  the  corresponding  one  with  hydrochloric  acid.  This  is 
a  further  illustration  of  the  fact  already  mentioned,  that  there  is  much 
less  affinity  between  hydrogen  and  bromine  than  between  hydrogen 
and  chlorine. 


158     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

IODINE 

Atomic  weight,  126.7.     Molecular  weight  (at  600°),  253.4. 

History  —  Iodine  was  discovered  by  Courtois  (1812)  in  the  course 
of  experiments  designed  to  prepare  potassium  nitrate  from  solutions 
obtained  by  lixiviating  the  ashes  of  sea-weeds.  The  name  is  derived 
from  ioetdri<r,  violet-coloured,  in  allusion  to  the  characteristic  colour  of 
iodine  vapour,  by  means  of  which  it  was  first  discovered. 

Occurrence  —  Iodine  does  not  occur  free  in  nature,  but  is  found 
chiefly  in  combination  with  potassium,  magnesium,  and  calcium  as 
iodides  and  iodates.  It  is  found  in  small  amount  in  sea-  water,  and 
also  in  certain  mineral  waters.  Certain  sea-weeds  have  the  power  of 
removing  iodine  from  sea-water,  and  these  sea-weeds  constitute  one 
of  the  chief  commercial  sources  for  this  element.  The  other  impor- 
tant source  is  Chili  saltpetre,  NaNO3  (p.  407),  with  which  it  is 
associated,  mainly  as  sodium  iodate,  NaIO3,  to  the  extent  of  about 
0.2  per  cent. 

Preparation  —  Iodine  may  be  obtained  from  iodides  by  either  of 
the  chief  methods  described  under  bromine  :  (a)  by  the  action  of 
chlorine  on  an  iodide: 

2KI  +  C12->2KC1  +  I2; 

(b)  by  the  action  of  a  mixture  of  sulphuric   acid   and  manganese 
dioxide  on  the  iodide  : 


If  the  first  method  is  used,  care  must  be  taken  not  to  use  excess  of 
chlorine,  which  leads  to  the  formation  of  a  compound  of  iodine  and 
chlorine  (p.  189),  and  consequent  loss  of  iodine.  The  commercial 
preparation  of  iodine  consists  in  obtaining  the  iodide  from  some 
natural  source,  and  then  liberating  the  iodine  by  one  of  the  above 
methods. 

(i)  From  Sea-weed  —  The  sea-  weed  is  first  dried  and  then 
either  burned  at  as  low  a  temperature  as  possible,  or,  better,  sub- 
jected to  dry  distillation  in  iron  retorts  at  a  low  red  heat.  The  ash 
(known  as  kelp  in  Scotland  and  varec  in  Normandy)  contains  practi- 
cally all  the  iodine  as  iodides,  as  well  as  chlorides  and  sulphates  ;  it 
is  extracted  with  water,  the  solution  evaporated,  and  the  potassium 
and  sodium  sulphates  and  chlorides  separated  as  far  as  possible  by 
crystallization.  The  mother  liquid  is  then  treated  with  sulphuric  acid 


THE    HALOGENS   AND   HALOGEN   ACIDS      159 

and  manganese  dioxide  and  gently  heated,  when  the  iodine  distils 
over  and  is  collected  in  a  series  of  earthenware  condensers. 

The  preparation  of  iodine  by  this  method  once  formed  an  important 
industry  on  the  northern  and  western  coasts  of  Scotland,  but  much  of 
the  iodine  used  in  commerce  is  obtained  from  Chili  saltpetre,  by  a 
method  now  to  be  described. 

(2)  As  already  mentioned,  Chili  saltpetre  (crude  sodium  nitrate) 
contains  a  small  proportion  of  sodium  iodate,  NaIO3.  As  the  latter 
contains  a  large  proportion  of  oxygen,  it  is  evident  that  it  must  be 
reduced  in  order  to  obtain  free  iodine.  The  reduction  can  be  effected 
by  adding  a  mixture  of  the  normal  sulphite  of  sodium,  Na2SO3,  and 
the  acid  sulphite,  NaHSO3.  The  iodine  is  precipitated  according  to 
the  equation 


It  is  freed  from  adhering  liquor  by  pressure  and  purified  by  sublima- 
tion. In  the  process  the  sulphites  are  oxidized  to  sulphate. 

Iodine  may  be  purified  by  mixing  with  a  little  powdered  potassium 
iodide  and  resubliming. 

Physical  Properties—  Iodine  occurs  in  the  form  of  large,  bluish- 
black,  lustrous  plates  of  density  4.95;  it  melts  at  114°  and/  boils  at 
184°.  The  vapour  is  reddish-violet  at  low  temperatures,  but  on  further 
heating  becomes  deep  blue. 

At  temperatures  up  to  500°  the  vapour  density  of  iodine  at  atmos- 
pheric pressure  is  about  127,  corresponding  with  the  formula  I2,  but 
as  the  temperature  is  further  increased  it  steadily  diminishes  and  at 
1500°  has  fallen  to  half  the  former  value.  As  in  the  case  of  bromine, 
we  must  assume  that  iodine  dissociates  on  raising  the  temperature, 
according  to  the  equation 


and  that  the  reaction  in  the  direction  of  the  upper  arrow  is  complete 
at  1500°  under  760  mm.  pressure.  Under  similar  conditions,  the 
dissociation  of  bromine  is  much  smaller.  Iodine  is  only  very  slightly 
soluble  in  water  (0.34  grams  per  litre  at  25°),  but  readily  dissolves  with 
a  violet  colour  in  carbon  disulphide,  chloroform,  and  certain  other 
organic  solvents,  and  with  a  brown  colour  in  alcohol,  ether,  and  in 
an  aqueous  solution  of  potassium  iodide.  This  difference  in  the  colour 
of  the  solutions  has  given  rise  to  considerable  discussion,  but  it  is  now 
generally  assumed  that  in  the  brown  solutions  iodine  enters  into 
combination  with  the  solvent,  whereas  there  is  little  or  no  combina- 


160     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

tion  in  the  violet  solutions.  This  view  is  supported  by  the  observation 
that  all  the  brown  solutions  become  violet  on  heating  ;  in  this  as  in 
other  cases  we  may  assume  that  the  compounds  gradually  decompose 
as  the  temperature  is  raised. 

When  the  brown  aqueous  solution  is  shaken  with  carbon  disulphide 
or  chloroform  it  passes  almost  entirely  into  the  organic  solvent,  which 
is  coloured  violet.  We  have  already  learnt  that  a  gas  in  contact  with 
a  solvent  may  be  regarded  as  being  distributed  between  the  space  and 
the  solvent,  the  ratio  of  the  distribution  depending  on  the  nature  of 
the  solvent.  The  phenomenon  now  under  consideration  is  closely 
analogous,  inasmuch  as  the  iodine  distributes  itself  between  the 
water  and  the  chloroform  in  a  ratio  depending  on  its  solubility  in  the 
two  solvents.  It  has  further  been  found  that  the  ratio  is  independent 
of  the  concentration  of  the  iodine  ;  in  other  words,  the  distribution 
follows  Henry's  Law  (p.  79). 

Chemical  Properties—  Iodine  acts  chemically  much  like 
chlorine  and  bromine,  but  usually  less  energetically  ;  it  combines 
directly  with  certain  non-metals  (such  as  phosphorus)  and  metals  to 
form  iodides.  It  combines  very  slowly  with  hydrogen,  and  even  after 
prolonged  heating  of  the  elements  at  high  temperatures  the  reaction 
is  incomplete.  Further,  if  hydrogen  iodide  is  heated  at  a  high 
temperature,  it  partly  breaks  down  into  hydrogen  and  iodine,  so  that 
we  are  dealing  with  a  reversible  reaction  represented  by  the  equation 


The  investigation  of  this  reaction  has  been  of  considerable  im- 
portance in  the  development  of  chemistry,  and  it  will  be  further 
referred  to  at  a  later  stage  (p.  164). 

Iodine  in  water,  like  chlorine,  is  an  oxidizing  agent,  and  the 
reaction  depends  upon  the  formation  of  hydriodic  acid  and  the 
liberation  of  oxygen  : 

I2+H2O->2HI  +  O. 

As,  however,  the  affinity  between  hydrogen  and  iodine  is  com- 
paratively small,  the  reaction  does  not  take  place  at  all  unless  a 
reducing  agent  (e.g.  sulphurous  acid,  p.  300),  is  present  to  remove 
the  oxygen.  This  may  be  taken  as  a  further  illustration  of  the 
effect  of"  the  heat  of  reaction  on  the  direction  of  a  chemical  change  ; 
in  virtue  of  the  additional  heat  given  out  by  the  oxidation  of 
sulphurous  acid,  the  total  change  becomes  strongly  exothermic  and 
proceeds  spontaneously. 


THE    HALOGENS   AND    HALOGEN    ACIDS        161 

The  most  characteristic  test  for  iodine  is  the  deep  blue  colour 
produced  when  even  traces  of  it  are  added  to  starch  solution. 
The  blue  compound  is  termed  "iodide  of  starch,"  but  its  nature 
is  not  well  understood  and  there  is  no  evidence  that  it  is  a  true 
chemical  compound.  On  warming  to  about  80°  the  colour  is 
discharged,  but  it  returns  as  the  solution  cools.  It  is  a  remark- 
able fact  that  the  colour  is  not  given  by  iodine  alone  in  water, 
but  only  when  a  soluble  iodide  is  also  present. 


HYDROGEN  IODIDE  (HYDRIODIC  ACID),  HI 

Preparation—  The  methods  used  in  preparing  hydrogen  iodide 
are  almost  identical  with  those  used  for  preparing  hydrogen  bromide, 
and  may  therefore  be  very  briefly  dealt  with. 

(1)  A   certain   amount   of  hydtpgen   iodide   is   obtained  when   a 
mixture   of   hydrogen   and    iodine   vapour    is  passed   over  heated, 
finely  divided  platinum,  but  under  all  circumstances  the  combina- 
tion is  only  partial. 

(2)  As  in  the  case   of  hydrogen   bromide,  hydrogen   iodide  can 
be    obtained   by   heating  an  alkali   iodide   with    phosphoric    acid  ; 
the  hydrogen   iodide,  being  volatile,  readily  passes   off.     Sulphuric 
acid   cannot   be   used    in    place   of   phosphoric    acid   because   the 
hydriodic  acid  first  formed  reduces  the  sulphuric  acid  ultimately  to 
hydrogen  sulphide,  thus, 

(i.)  2KI+H2SO4=K2SO4+2HI. 
(ii.)  H2S 


As  hydrogen  iodide  is  a  much  more  powerful  reducing  agent  than 
hydrogen  bromide,  very  little  hydrogen  iodide  is  given  off  in  the 
above  process. 

(3)  An  aqueous  solution  of  hydrogen  iodide  is  readily  obtained 
by  passing  hydrogen  sulphide  through  iodine  suspended  in  water. 
The  reaction  is  a  reversible  one,  as  represented  by  the  equation 


and  is  incomplete  if  a  fairly  high  concentration  of  hydriodic  acid 
is  reached.  In  dilute  solution,  however,  the  reaction  proceeds 
practically  to  completion  in  the  direction  of  the  upper  arrow  (cf. 
p.  148).  As  already  mentioned,  the  total  change  is  exothermic  in 
virtue  of  the  great  heat  of  solution  of  hydriodic  acid. 

(4)  The  most  satisfactory  method  of  preparing  hydrogen  iodide 
II 


1 62     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


is  to  mix  red  phosphorus  and  iodine  in  a  dry  flask  and  allow  water 
to  drop  slowly  on  the  mixture  (Fig.  44). 


The  escaping  gas  is  freed  from  traces  of  free  iodine  by  passing 

it  over  red  phosphorus  in 
a  U-tube,  B,  and  may  be 
collected  by  upward  dis- 
placement of  air  or  over 
mercury  (p.  39).  If  an  aque- 
ous solution  is  required,  the 
gas  is  delivered  through  a 
tube  of  the  form  shown, 
which  dips  just  below  the 
surface  of  the  water.  Should 
the  water  be  sucked  back 
during  the  operation,  it  is 
prevented  from  entering  the 
U-tube  by  the  wide  bulb  on 
the  delivery  tube. 

Physical  Properties 
—  Hydrogen  iodide  is  a 
colourless  gas  with  a  sharp 
odour  ;  it  fumes  in  contact 
with  moist  air  (p.  163).  It 
can  readily  be  condensed 
to  a  colourless  liquid  which 
boils  at  -35.  7°  and  melts 
at  -  50.8°.  One  volume  of 
water  dissolves  425  volumes 
of  the  gas  at  10°,  forming 
a  strongly  acid  solution 
(hydriodic  acid).  When 
distilled,  hydriodic  acid,  like 
the  other  two  hydrogen 
acids  already  referred  to, 


FIG.  44. 


forms  a  mixture  of  constant  boiling-point  (127°  at  760  mm.  pressure) 
containing  57.7  per  cent,  of  hydrogen  iodide. 

Hydrogen  iodide  is  an  endothermic  compound  ;  the  thermochemical 
equation  for  its  formation  in  the  gaseous  form  is 

-2  x6,ioo  cal. 


THE   HALOGENS   AND    HALOGEN   ACIDS        163 

The  heat  of  solution  of  hydrogen  iodide  in  water  is  high,  amount- 
ing to  19,200  cal.  per  mol. 

Chemical  Properties  —  The  aqueous  solution  of  hydrogen  iodide 
is  a  typical  acid,  which  dissolves  many  metals  to  form  salts,  the  iodides. 
The  latter  resemble  the  chlorides  and  bromides  in  chemical  properties, 
most  of  them  are  soluble  in  water.  The  chief  differences  in 
behaviour  between  hydriodic  acid  and  the  other  halogen  acids  are 
connected  with  the  comparatively  small  affinity  between  hydrogen 
and  iodine,  already  referred  to.  For  this  "reason  hydriodic  acid 
is  a  powerful  reducing  agent,  hydrogen  being  given  up  in  the 
process  and  iodine  set  free.  An  illustration  of  this  behaviour  has 
already  been  met  with  in  the  reaction  between  sulphuric  acid  and 
potassium  iodide  (p.  161).  Hydriodic  acid  is  largely  used  in  organic 
chemistry  for  reducing  purposes. 

In  the  same  connexion  is  the  fact  that  solutions  of  hydriodic  acid 
rapidly  turn  brown  in  the  air  owing  to  liberation  of  iodine  : 


Although  iodine  is  practically  insoluble  in  water,  it  readily  dissolves 
in  the  excess  of  hydriodic  acid. 

It  has  already  been  mentioned  that  hydrogen  iodide  splits  up 
partially  on  heating  according  to  the  reversible  reaction 


The  amount  of  decomposition  is  the  greater  the  higher  the 
temperature. 

Salts  of  the  Halogen  Acids—  As  the  halogen  acids  are  mono- 
basic, their  salts  with  univalent  metals  are  of  the  type  MX  [M  =  metal, 
X  =  halogen],  with  divalent  metals  of  the  type  MX2,  with  trivalent 
metals  of  the  type  MX3,  and  so  on.  In  accordance  with  the  general 
methods  of  preparing  salts  (p.  99)  they  can  be  obtained  by  action  of 
the  appropriate  acid  on  the  metal,  oxide,  or  hydroxide,  and  in  certain 
cases  by  direct  combination  of  metal  and  halogen.  The  following 
equations  illustrate  the  different  methods  of  salt  formation  :  — 

Zn  +  2HBr=ZnBr2  +  H2 
ZnO  +  2HBr=ZnBr2  +  H2O 


CHAPTER  XIII 
CHEMICAL    EQUILIBRIUM—  THERMAL    DISSOCIATION 

IN  the  present  section  the  general  principles  underlying  certain 
changes  already  referred  to  will  be  more  fully  considered.  We 
have  seen  that  when  a  definite  amount  of  hydrogen  iodide  is  sealed 
up  in  a  glass  tube  and  heated  in  the  vapour  of  boiling  sulphur  it 
begins  to  split  up  into  hydrogen  and  iodine,  but  the  change  stops 
when  about  2  1  per  cent,  of  it  is  decomposed.  Under  these  circum- 
stances, therefore,  the  mixture  contains  79  per  cent,  of  hydrogen 
iodide  and  21  per  cent,  of  hydrogen  and  iodine,  and  the  proportion 
remains  unaltered  at  445°,  no  matter  how  long  the  mixture  is  heated. 
On  the  other  hand,  if  equivalent  amounts  of  hydrogen  and  iodine, 
contained  in  a  sealed  tube,  are  heated  at  445°  till  no  further  change 
occurs,  it  is  found  that  the  resulting  mixture  of  gases  contains  79  per 
cent,  of  hydrogen  iodide  and  21  per  cent,  of  hydrogen  and  iodine. 
It  is  evident,  therefore,  that  just  as  water  is  in  equilibrium  with  a 
certain  concentration  of  water  vapour  at  a  definite  temperature,  so 
hydrogen,  iodine  and  hydrogen  iodide  are  in  equilibrium  at  a  definite 
temperature  when  a  certain  concentration  of  each  is  present.  The 
equilibrium  between  water  and  water  vapour  is  often  termed  a  physical 
equilibrium,  whilst  that  between  hydrogen,  iodine  and  hydrogen 
iodide,  being  reached  as  the  result  of  a  chemical  change,  is  called 
a  chemical  equilibrium.  All  the  facts  are  conveniently  represented 
by  the  equation  for  a  reversible  reaction  — 


which  has  already  been  fully  explained. 

The  kinetic  theory  throws  a  great  deal  of  light  on  the  nature  of 
equilibria  of  this  kind.  As  the  mere  presence  of  hydrogen  and  iodine 
does  not  retard  the  decomposition  of  hydrogen  iodide,  it  may  appear 
at  first  sight  as  if  complete  decomposition  of  the  latter  ought  to  occur 
if  sufficient  time  be  allowed.  As,  however,  hydrogen  and  iodine  can 
combine  to  form  hydrogen  iodide  under  the  same  conditions,  this 

reaction  sets  in  as  soon  as  any  of  the  elements  are  present,  and  the 

164 


CHEMICAL   EQUILIBRIUM  165 

apparent  static  equilibrium  is  really  a  kinetic  one,  being  the  point  at 
which  the  amount  of  hydrogen  iodide  decomposed  in  a  given  time  is 
just  balanced  by  the  amount  formed  by  recombination  of  its  com- 
ponent elements. 

We  are  now  in  a  position  to  consider  rather  more  fully  the 
influence  of  the  conditions  upon  the  equilibrium  in  a  chemical  system. 
The  first  point  to  notice  is  that  an  equilibrium  can  only  be  established 
when  all  the  reacting  substances  remain  in  the  system.  This  follows 
at  once  from  the  kinetic  interpretation  of  a  chemical  equilibrium  —  if 
one  of  the  products  of  the  reaction  is  continually  removed  from  the 
system  the  back  reaction  is  no  longer  possible.  When  excess  of 
sulphuric  acid  is  added  to  sodium  chloride  at  the  ordinary  tempera- 
ture, the  equilibrium  in  the  solution  is  represented  by  the  equation 

NaCl  +  H2SO4 


the  amount  of  the  products  formed  in  a  given  time,  as  represented  by 
the  upper  arrow,  being  just  balanced  by  the  amount  of  the  original 
substances  reformed  by  interaction  of  the  products.  If,  however,  the 
mixture  is  heated,  the  hydrogen  chloride,  being  readily  volatile, 
escapes  from  the  system  as  fast  as  it  is  formed,  the  reaction  repre- 
sented by  the  lower  arrow  cannot  take  place,  and  therefore  the 
chemical  change  proceeds  practically  to  completion  in  the  direction 
of  the  upper  arrow.  It  should  be  carefully  noted  that  sulphuric  acid 
does  not  displace  hydrochloric  acid  from  combination  with  sodium 
because  the  former  is  the  stronger  acid  ;  in  fact,  as  we  shall  learn 
later,  hydrochloric  acid  is  stronger  than  sulphuric  acid  under  equiva- 
lent conditions. 

Similar  considerations  account  for  the  fact  that  the  action  of  sulphuric 
acid  on  barium  peroxide,  represented  by  the  equation  (p.  138) 

Ba02  +  H2S04^BaS04  j  +H2O2, 

proceeds  almost  to  completion  in  the  direction  of  the  upper  arrow. 
As  barium  sulphate  is  practically  insoluble  in  water,  it  is  removed 
from  the  system  in  the  solid  form  as  fast  as  it  is  produced,  so  that 
the  back  reaction,  represented  by  the  lower  arrow,  does  not  take 
place  to  any  appreciable  extent. 

Influence  of  Concentration  of  Reacting  Substances 
on  Equilibrium.  Law  of  Mass  Action  —  The  next  point  to 
be  considered  is  the  effect  on  the  equilibrium  of  varying  the  relative 
amounts  of  the  reacting  substances.  Suppose,  for  instance,  we  repeat 
the  experiment  of  heating  together  hydrogen  and  iodine,  with  the 


166     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

single  difference  that  double  the  quantity  of  iodine  is  taken  for  the 
same  volume.  It  is  then  found  that,  instead  of  79  per  cent.,  about 
93  per  cent,  of  the  total  quantity  of  hydrogen  is  converted  into 
hydrogen  iodide  at  equilibrium  ;  in  other  words,  with  regard  to  the 
reaction 


an  increase  in  the  proportion  of  iodine  displaces  the  equilibrium  in 
the  direction  of  the  upper  arrow.  If  instead  of  adding  more  iodine 
the  proportion  of  hydrogen  is  increased,  the  effect  on  the  equilibrium 
is  in  the  same  direction. 

From  the  kinetic  standpoint,  this  means  that  an  increase  in  the 
proportion  of  hydrogen  or  of  iodine  in  the  system  causes  a  relative 
increase  in  the  rate  of  the  reaction  represented  by  the  upper  arrow. 
According  to  the  molecular-kinetic  theory,  the  molecules  of  hydrogen 
and  iodine  must  come  into  contact  in  order  that  hydrogen  iodide  may 
be  formed,  and  it  is  plausible  to  suppose  that  the  rate  of  formation 
of  the  iodide  is  proportional  to  the  number  of  collisions  per  unit  time 
between  the  reacting  molecules.  If  in  a  definite  volume  of  the  gases 
the  amount  of  iodine  is  doubled,  the  number  of  collisions  per  unit 
time  will  also  be  approximately  doubled,  and  the  rate  of  the  reaction 
represented  by  the  upper  arrow  correspondingly  increased.  A  new 
state  of  equilibrium  will  finally  be  reached  when  the  amounts  of 
hydrogen  iodide  formed  and  decomposed  in  unit  time  again  balance, 
and  this  is  only  possible  by  a  falling  off  in  the  amounts  of  hydrogen 
and  of  iodine  in  the  given  volume  and  a  corresponding  increase  in 
the  amount  of  hydrogen  iodide.  This  means  that  the  equilibrium  is 
displaced  in  the  direction  of  the  upper  arrow  by  the  addition  of 
hydrogen  or  of  iodine,  in  accordance  with  the  experimental  result. 

From  the  experimental  data  of  this  and  similar  experiments,  the 
exact  relationship  between  the  rate  of  a  chemical  reaction  and  the 
amounts  of  the  reacting  substances  present  can  be  calculated.  It 
would  lead  too  far  to  work  the  matter  out  in  detail,1  but  the  result 
may  be  briefly  stated  as  follows  :  The  rate  of  a  chemical  reaction  is 
Proportional  to  the  molecular  concentration  of  each  of  the  reacting 
substances.  It  is  important  to  note  that  the  rate  is  not  proportional 
to  the  amount  of  each  substance  present,  but  to  its  concentration, 
this  is,  to  the  amount  per  unit  volume.  This  result  is  also  in  entire 
accord  with  the  above  molecular-kinetic  considerations.  If  two  vessels 
of  equal  volume  contain  the  same  amount  of  iodine  and  quantities  of 

1  Cf.  Physical  Chemistry,  p.  156. 


CHEMICAL   EQUILIBRIUM  167 

hydrogen  in  the  ratio  i :  2,  the  number  of  collisions  and  therefore  the 
speed  of  the  reaction  in  the  second  vessel  will  be  double  that  in  the 
first  vessel,  corresponding  with  the  fact  that  the  concentration  of 
hydrogen  in  the  former  is  double  that  in  the  latter.  If,  however,  the 
contents  of  the  first  tube  are  mixed  with  an  equivalent  amount  of 
hydrogen,  so  that  the  resulting  volume  is  greater  than  at  first,  the 
quantity  of  hydrogen  is  doubled,  but  not  its  concentration,  and  it  is 
an  experimental  fact  that  the  speed  of  the  reaction  is  not  doubled. 
It  is  evident  that  owing  to  the  greater  free  space  which  the  molecules 
occupy  in  the  latter  case  the  number  of  collisions  is  not  doubled  by 
doubling  the  amount  of  hydrogen. 

The  convenience  of  expressing  concentrations  in  mols  per  litre 
instead  of  in  parts  by  weight  will  be  obvious  from  what  has  been 
stated  in  previous  chapters.  As  chemical  reactions  between  different 
substances  always  take  place  in  molar  or  molecular  proportions,  there 
is  a  great  gain  in  simplicity. 

The  important  result  just  stated,  that  the  rate  of  a  chemical  change 
(in  other  words,  the  amount  of  a  chemical  change  taking  place  in  a 
given  time)  is  proportional  to  the  molecular  concentration  of  each  of 
the  reacting  substances,  is  usually  called  the  Law  of  Mass  Action. 
Instead  of  "molecular  concentration"  the  term  "active  mass"  is 
often  used  in  stating  the  law  ;  but  the  former  method  is  preferable, 
as  the  latter  statement  appears  to  imply  that  the  rate  of  a  chemical 
change  is  doubled  by  doubling  the  mass  of  one  of  the  reacting  sub- 
stances, which  is  not  the  case  unless  the  total  volume  is  kept  constant. 

The  substance  of  the  present  section  may  be  partially  summarized 
by  a  further  consideration  of  the  reaction  represented  by  the  equation 

NaCl+H2SO4^NaHSO4+HCl. 

Assuming  that  under  the  conditions  of  the  experiment  all  the  reacting 
substances  remain  in  the  system,  the  equilibrium  is  determined  by 
two  distinct  kinds  of  factor : — 

(1)  The  specific  chemical  affinity  between  the  sulphuric  acid  and 
sodium  chloride  and  that  between  the  hydrochloric  acid  and  sodium 
hydrogen  sulphate.     These  specific  affinities  are  assumed  to  be  inde- 
pendent of  the  concentration,  but  depend  on  other  factors,  such  as 
temperature,  etc. 

(2)  The  relative  concentrations  v  of  the  reacting  substances.     The 
relationship  between   the   molar   concentration   and   the  amount  of 
chemical  action  is  expressed  in  the  law  of  mass  action,  already  fully 
discussed. 


1  68     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

This  method  of  representing  the  facts,  although  not  in  all  respects 
satisfactory,  will  be  found  suitable  for  our  present  purpose. 

These  considerations  can  be  put  much  more  clearly  in  mathematical  form.1 
Assume  that  two  substances  A  and  B  react  to  form  two  new  substances  C  and 
D,  according  to  the  equation 


and  that  the  respective  molar  concentrations  at  equilibrium  are  a,  b,  c,  and  d. 
It  is  clear  from  the  foregoing  that  the  rate  of  the  direct  reaction  is  proportional 
both  to  a  and  to  6,  and  is  therefore  proportional  to  their  product.  We  may  put 
this  in  the  following  form  — 

Rate  direct  =kab, 

where  k  is  a  constant  depending  on  the  affinity  between  A  and  B.  In  the  same 
way 

Rate  re  verse  =Kcd, 

where  k'  is  a  second  constant,  which  depends  on  the  affinity  between  C  and  D. 
At  equilibrium  the  rates  of  the  direct  and  reverse  actions  are  equal  by  definition, 
that  is, 

kab^k'cd. 

The  above  equation  may  be  written 

-A=K 

cd    k 

where  K,  being  the  ratio  of  the  two  constants  k'  and  k,  is  also  constant,  and  like 
k  and  k  ',  is  independent  of  the  concentrations.  K  is  often  termed  the  equilibrium 
constant. 

The  above  result  may  be  put  in  the  following  form  :  At  equilibrium  the  product 
of  the  concentrations  on  one  side,  divided  by  the  product  of  the  concentrations  on  the 
other  side,  is  constant  at  constant  temperature.  This  is  the  most  general  statement 
of  the  Law  of  Mass  Action. 

When  a  molecule  each  of  A  and  B  react  to  form  a  molecule  of  a  third  substance 
C,  according  to  the  equation 

A+B^tC, 
the  equilibrium  equation  ab—Kcd  simplifies  to  ab=^Kc. 

The  law  of  mass  action  can  be  experimentally  illustrated  by  means 
of  the  reversible  reaction  between  ferric  chloride,  FeCl3,  and  ammo- 
nium thiocyanate,  NH4CNS,  which  form  ferric  thiocyanate,  Fe(CNS)3, 
and  ammonium  chloride,  NH4C1,  according  to  the  equation 


Solutions  of  the  salts  are  first  prepared  ;  the  thiocyanate  solution  con- 
tains 3.7  grams  of  the  salt  to  100  c.c.  of  water  and  the  ferric  chloride 

1  For  details  see  Physical  Chemistry,  chap.  vii. 


THERMAL   DISSOCIATION  169 

solution  3  grams  of  the  commercial  salt  and  12.5  c.c.  of  concen- 
trated hydrochloric  acid  to  100  c.c.  of  water.  Five  c.c.  of  each  of  the 
solutions  are  added  to  2  litres  of  water,  and  the  solution  divided 
between  four  beakers.  Of  the  four  salts  present  in  solution,  the 
ammonium  salts  are  colourless,  ferric  chloride  is  very  pale  red  in 
dilute  solution,  and  ferric  thiocyanate  deep  blood-red.  As  the  con- 
tents of  the  four  beakers  prepared  as  above  are  pale-red,  it  is  evident 
that  the  equilibrium  is  considerably  displaced  in  the  direction  of  the 
lower  arrow.  To  the  contents  of  two  of  the  beakers  are  added  5  c.c. 
of  the  ferric  chloride  and  thiocyanate  solutions  respectively,  and  it 
will  be  observed  that  the  solutions  become  blood-red,  owing  to  the 
displacement  of  the  equilibrium  in  the  direction  of  the  upper  arrow, 
in  accordance  with  the  law  of  mass  action.  On  the  other  hand,  the 
addition  of  50  c.c.  of  a  concentrated  solution  of  ammonium  chloride 
to  the  contents  of  the  third  beaker  makes  it  practically  colourless  ; 
the  equilibrium  is  thus  displaced  in  the  direction  of  the  lower  arrow, 
again  in  accordance  with  the  law  of  mass  action.  The  solution  in 
the  fourth  beaker  is  kept  for  comparison. 

Thermal  Dissociation—  When  a  chemical  compound  splits 
up  into  simpler  substances  on  heating,  and  the  products  are  capable 
of  recombining  to  form  the  original  compound  on  cooling,  the  latter 
is  said  to  undergo  dissociation.  It  is  clear  that  the  term  only  applies 
to  reversible  reactions  ;  the  decomposition  of  potassium  chlorate  by 
heat  into  potassium  chloride  and  oxygen,  an  irreversible  reaction,  is 
not  a  dissociation.  A  number  of  instances  of  dissociation  by  heat 
have  already  been  met  with,  such  as  the  splitting  up  of  hydrogen  iodide 
into  its  elements  (p.  164),  the  splitting  up  of  steam,  according  to  the 
equation  2H2O^I2H2  +  O2;  of  iodine,  according  to  the  equation 
I2<^2l,  and  so  on.  Many  other  illustrations  of  this  process  will  be 
met  with  later.  Owing  to  its  importance  a  further  example  will  be 
mentioned  here,  namely,  the  dissociation  of  phosphorus  pentachloride, 
PC16,  into  the  trichloride,  PC13,  and  free  chlorine  ; 


The  methods  employed  in  detecting  and  measuring  thermal  dis- 
sociation depend  on  the  nature  of  the  reaction.  The  dissociation  of 
hydrogen  iodide  at  high  temperatures  can  be  measured  by  cooling 
the  mixture  rapidly  to  room  temperature,  which  stops  the  re- 
action, and  determining  the  relative  concentrations  of  the  reacting 
substances  at  leisure.  A  modification  of  this  method,  which  may  be 
used  to  show  the  occurrence  of  dissociation  in  water  vapour  at  high 


170     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

temperatures,  has  already  been  described  (p.  32).  In  the  case  of 
phosphorus  pentachloride,  the  latter  and  the  bichloride  are  nearly 
colourless  in  the  form  of  vapour,  whilst  chlorine  is  light  green,  and 
the  progress  of  dissociation  as  the  temperature  is  raised  is  indicated 
by  the  gradual  increase  in  the  depth  of  colour.  The  phenomenon 
can  be  illustrated  still  more  satisfactorily  by  using  phosphorus  penta- 
bromide,  PBr5,  which  dissociates  in  an  analogous  way,  owing  to  the 
pronounced  colour  of  the  bromine. 

From  a  quantitative  point  of  view  the  progress  of  dissociation  can 
often  be  followed  by  measurements  of  density.  When  iodine  exists 
as  I2  molecules  its  molecular  weight  is  254,  and  its  density,  referred 
to  hydrogen  at  the  same  temperature,  is  127.  If,  however,  it  is  com- 
pletely split  up  into  I  atoms  its  molecular  weight  (which  under  these 
circumstances  is  the  same  as  the  atomic  weight)  is  127,  and  its  density 
63.5.  Victor  Meyer  found  that  its  density  at  1250°  is  about  81.5,  a 
number  which  lies  between  that  for  I2  molecules  and  I  atoms.  We 
therefore  assume  that  the  element  has  undergone  partial  dissociation 
according  to  the  equation 


and  it  can  readily  be  calculated  that  the  density  at  1250°  corresponds 
with  a  degree  of  dissociation  of  about  72  per  cent. 

The  above  remarks  on  the  experimental  investigation  of  dissociation 
apply  also  to  chemical  equilibria,  of  which  dissociation  is  simply  a 
special  type. 

An  important  question  in  connexion  with  the  present  subject  is  the 
effect  of  excess  of  one  of  the  products  of  dissociation  of  a  substance 
on  its  degree  of  dissociation.  When  a  definite  weight  of  phosphorus 
pentabromide  is  vaporized  in  a  closed  vessel  at  say  200°,  the  vapour 
is  deep  red,  indicating  the  presence  of  a  considerable  proportion  of 
free  bromine  : 


If,  however,  the  vessel  previously  contains  excess  of  the  tribromide,  it. 
•will  be  found  that  the  vapour  is  only  very  slightly  coloured,  indicating 
that  under  conditions  otherwise  the  same  the  degree  of  dissociation 
is  much  less  in  the  presence  of  excess  of  one  of  the  products  of  dis- 
sociation. The  same  fact  can  also  be  shown  quantitatively  by 
determining  the  density  of  the  vapour  of  the  pentabromide  alone  and 
in  the  presence  of  excess  of  the  tribromide.  It  is  considerably  highei 
in  the  latter  case,  indicating  less  complete  dissociation. 


THERMAL   DISSOCIATION  171 

The  above  conclusion  also  follows  at  once  from  the  kinetic  inter- 
pretation of  the  law  of  mass  action.  The  addition  of  more  tribromide 
increases  the  reaction  velocity  in  the  direction  of  the  lower  arrow, 
but  does  not  affect  the  specific  rate  at  which  the  pentabromide 
decomposes  (represented  by  the  upper  arrow),  so  that  the  equilibrium 
is  displaced  towards  the  left ;  in  other  words,  the  degree  of  dissocia- 
tion is  diminished.  This  result  may  be  stated  as  follows :  The  degree 
of  dissociation  of  a  compound  is  diminished  by  addition  of  one  of  the 
Products  of  dissociation^  provided  that  the  volume  is  kept  constant. 

Homogeneous  and  Heterogeneous  Equilibria— So  far 
attention  has  been  confined  almost  entirely  to  equilibrium  in  homo- 
geneous systems — systems  which  consist  of  a  single  form  of  matter 
and  are  of  the  same  composition  at  all  points.  For  the  sake  of  com- 
pleteness, reference  must  also  be  made  to  heterogeneous  systems, 
which  consist  of  two  or  more  distinct  portions — so-called  phases — 
which  are  themselves  of  uniform  composition  but  are  separated  by 
definite  surfaces  from  other  phases  (p.  69).  Liquid  water  in  equi- 
librium with  its  vapour  is  a  heterogeneous  system  made  up  of  two 
phases,  water  and  vapour,  but  the  equilibrium  in  this  case  is  of  a 
physical  nature.  A  more  complicated  heterogeneous  system  is  that 
in  a  saturated  solution  of  a  salt ;  in  this  case  there  are  three  phases, 
namely  solid  salt,  solution  and  vapour  in  equilibrium.  A  more 
detailed  consideration  of  heterogeneous  systems  will  be  given  at  a 
later  stage. 

Effect  of  Change  of  Temperature  and  Pressure  on 
Chemical  Equilibrium.  Le  Chatelier's  Theorem— When 
the  temperature  of  a  system  in  equilibrium  is  altered,  two  main  types 
of  change  may  occur — (i)  the  equilibrium  may  be  displaced,  all  the 
original  substances  remaining  in  the  system ;  (2)  one  or  more  of  the 
phases  may  disappear,  an  equilibrium  of  an  entirely  different  type 
resulting.  To  take  an  example  of  the  latter  case  first,  if  a  mixture  of 
ice  and  water  in  equilibrium  with  water  vapour  (p.  68)  is  raised  in 
temperature  even  a  fraction  of  a  degree,  and  kept  at  the  new  tempera- 
ture, the  ice  melts  and  a  new  equilibrium  is  established  between 
water  and  water  vapour  (cf,  curve  OA,  p.  68).  In  the  same 
way,  when  certain  systems  in  chemical  equilibrium  at  a  definite 
temperature  are  heated  or  cooled  one  of  the  phases  may  entirely 
disappear. 

For  our  present  purpose  equilibria  of  the  first  type,  those  in  which 
alteration  of  temperature  produces  only  a  displacement  of  the  equi- 
librium, are  more  important.  It  has  been  found  that  the  direction  in 


172     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

which  the  equilibrium  is  displaced  is  closely  connected  with  the  heat 
given  out  or  absorbed  in  the  chemical  change.  In  order  to  simplify 
matters  as  much  as  possible,  we  assume  that  the  voUime  is  kept  con- 
stant throughout.  The  rule  connecting  heat  of  reaction  and  displace- 
ment of  equilibrium  may  now  be  stated  as  follows:  At  constant 
volume,  increase  of  temperature  favours  the  reaction  in  which  heat  is 
absorbed,  lowering  of  temperature  favours  the  contrary  reaction. 
This  rule  may  first  be  illustrated  by  means  of  the  ozone-oxygen 
equilibrium  (p.  135)  — 


As  heat  is  absorbed  in  the  formation  of  ozone,  a  rise  of  temperature 
ought  to  displace  the  equilibrium  towards  the  left  ;  in  other  words,  the 
proportion  of  ozone  in  equilibrium  with  oxygen  should  be  the  greater 
the  higher  the  temperature,  and  the  experimental  facts  are  in  entire 
accord  with  this  deduction.  The  corresponding  deduction,  that  at 
low  temperatures  the  equilibrium  is  very  near  the  right-hand  side, 
has  also  been  confirmed  experimentally. 

We  shall  meet  later  with  many  illustrations  of  this  rule.  One  very 
important  instance  has  already  been  mentioned,  namely,  the  effect  of 
change  of  temperature  on  the  solubility  (p.  83).  It  follows  at  once 
that  elevation  of  temperature  should  increase  the  solubility  of  salts 
such  as  potassium  nitrate,  which  dissolve  with  absorption  of  heat, 
and  diminish  the  solubility  of  compounds  such  as  calcium  hydroxide, 
which  dissolve  with  evolution  of  heat.  It  will  also  be  evident  why 
the  rule  only  applies  to  the  solubility  of  substances  in  solutions  already 
practically  saturated  with  them,  since  it  applies  only  to  systems  which 
are  approximately  in  equilibrium. 

If  no  heat  is  given  out  on  displacement  of  equilibrium,  the  latter 
ought  not  to  be  affected  by  altering  the  temperature  at  constant 
volume.  This  deduction  also  is  borne  out  by  experiment. 

The  position  of  equilibrium  is  also  generally  altered  by  change  of 
pressure  ;  the  rule  is  as  follows  :  At  constant  temperature  increase  of 
pressure  displaces  the  equilibrium  in  the  direction  in  which  the  volume 
diminishes,  whilst  decrease  of  pressure  has  the  contrary  effect.  As  an 
illustration  we  will  again  take  the  ozone-oxygen  equilibrium— 

203^30,5. 

Increase  of  pressure  should  displace  the  equilibrium  in  the  direction 
of  diminution  of  volume,  and  thus  favour  the  production  of  ozone, 
whereas  a  lowering  of  pressure  should  have  the  contrary  effect.  As 


THERMAL   DISSOCIATION  173 

it  happens,  the  rule  cannot  readily  be  tested  directly  in  this  Case 
on  account  of  experimental  difficulties,  but  it  has  been  proved 
valid  for  hundreds  of  other  equilibria  (some  of  which  will  be  men- 
tioned subsequently),  and  may  safely  be  assumed  to  hold  for  this 
case  also. 

It  follows  at  once  from  this  rule  that  if  the  volume  does  not  change 
on  displacement  of  equilibrium  at  constant  temperature,  alteration  of 
pressure  should  have  no  effect  on  the  position  of  equilibrium.  Such 
a  system  is  that  containing  hydrogen,  iodine  and  hydrogen  iodide : 

H2    +     I2  ^  2HI 

I  VOl.         I   VOl.       2  VOls. 

and  in  accordance  with  this  deduction  it  has  been  found  by  Bodenstein 
that  the  equilibrium  point  is  the  same,  within  the  limits  of  experi- 
mental error,  at  total  pressures  varying  between  wide  limits. 

As  the  rule  implies,  the  influence  of  change  of  pressure  on  equilibria 
in  liquid  and  solid  systems  is  very  slight,  corresponding  with  the 
small  changes  of  volume  in  such  systems. 

The  above  rules  with  regard  to  the  influence  of  temperature  and 
pressure  on  physical  and  chemical  equilibria  are  special  cases  of  a 
very  important  rule  or  theorem  usually  associated  with  the  name  of 
the  French  chemist  Le  Chatelier.  The  theorem  may  be  stated  as 
follows  : 

If  one  or  more  of  the  factors  determining  an  equilibrium,  namely,  con- 
centration, pressure  or  temperature,  is  altered,  the  equilibrium  becomes 
displaced  in  the  direction  which  tends  to  netttralize  the  effect  of  the 
alteration.  The  student  should  satisfy  himself  that  all  the  examples 
already  mentioned  in  this  section  are  in  accordance  with  the  rule, 
and  other  illustrations  of  it  will  readily  occur  to  the  mind.  For 
example,  the  conclusions  can  at  once  be  drawn  that  ice  must  be 
melted  by  raising  the  temperature  at  constant  pressure,  as  heat 
is  absorbed  in  the  process,  and  further,  increase  of  pressure  at  con- 
stant temperature  should  also  cause  ice  to  melt,  as  this  change  is 
attended  with  diminution  of  volume.  As  has  already  been  pointed 
out,  these  deductions  are  confirmed  by  experiment. 

Velocity  of  Reaction— We  have  already  met  with  a  number 
of  illustrations  of  the  fact  that  the  speed  of  chemical  changes  depend? 
greatly  on  the  nature  of  the  substances,  and  is  also  markedly  influenced 
by  the  conditions.  When  hydrochloric  acid  and  sodium  hydroxide 
are  mixed,  the  rate  of  reaction,  as  shown  by  the  change  of  colour  of 
the  indicator,  is  practically  instantaneous.  On  the  other  hand,  a 


174     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

mixture  of  hydrogen  and  oxygen  may  be  kept  for  years  at  the 
ordinary  temperature  without  any  apparent  change  taking  place 
(p.  37),  but  if  the  temperature  is  sufficiently  high  combination  at 
once  occurs. 

The  remarkable  influence  of  so-called  catalysts  in  accelerating 
chemical  changes  has  already  been  repeatedly  referred  to,  and  many 
other  illustrations  of  catalytic  action  will  be  met  with  in  the  course 
of  the  book. 

Light  has  a  remarkable  effect  in  accelerating  certain  chemical 
changes,  such  as  the  combination  of  hydrogen  and  chlorine,  but  it 
appears  to  have  little  or  no  influence  on  the  great  majority  of  chemical 
reactions. 

On  the  other  hand,  practically  all  chemical  changes  are  remarkably 
accelerated  by  rise  of  temperature.  A  careful  study  of  this  effect  has 
shown  that  the  speed  of  chemical  changes  is  generally  doubled  or 
trebled  for  a  rise  of  temperature  of  10°.  Suppose  we  assume  that 
the  speed  is  doubled  for  this  increment  of  temperature.  It  can 
readily  be  calculated  that  the  rate  of  reaction  is  increased  more 
than  1000  times  by  raising  the  temperature  by  100°,  and  more  than 
1,000,000  times  by  raising  the  temperature  by  200°.  It  is  therefore 
easy  to  understand  that  a  reaction  which  proceeds  fairly  rapidly  at 
a  high  temperature  may  be  so  slow  at  the  ordinary  temperature  that 
no  change  can  be  detected  in  weeks  or  even  months.  Instances 
of  this  which  have  already  been  met  with  are  the  combination  of 
hydrogen  and  chlorine  to  form  hydrogen  chloride,  and  of  hydrogen 
and  oxygen  to  form  water. 

We  are  now  in  a  position  to  understand  fully  the  phenomena 
associated  with  the  oxygen-ozone  equilibrium,  and  this  example  will 
serve  to  summarize  some  of  the  general  principles  just  discussed. 
When  a  silent  electric  discharge  is  passed  through  oxygen  a  mixture 
of  the  latter  with  ozone  is  obtained,  and  the  ozone  concentration  is 
much  higher  than  corresponds  with  the  thermal  equilibrium  at  the 
ordinary  temperature.  We  would  therefore  expect  that  as  soon  as 
the  mixture  is  removed  from  the  influence  of  the  electric  discharge 
the  ozone  will  begin  to  decompose,  and  doubtless  this  is  the  case. 
Owing,  however,  to  the  very  small  velocity  of  this  reaction  at  the 
ordinary  temperature,  the  mixture  appears  to  be  practically  stable. 
When,  however,  the  temperature  is  raised  to  250°,  the  speed  with 
which  the  system  moves  towards  equilibrium  is  enormously  in- 
creased and  the  ozone  rapidly  decomposes.  It  has  been  calculated 
that  the  concentration  of  ozone  in  true  equilibrium  with  oxygen 


THERMAL  DISSOCIATION  175 

at  atmospheric  pressure  and  ordinary  temperature  is  only  about 
o.ooi  per  cent. 

The  above  considerations  show  that  it  is  of  the  utmost  importance 
to  draw  a  clear  distinction  between  the  two  effects  of  temperature 
(i)  in  displacing  the  equilibrium  ;  (2)  in  accelerating  chemical 
changes. 

Allotropic  Modifications.  Isomerism.  Poly- 
morphism. Polymerism — When  the  same  element  exists  in 
different  forms,  the  forms  are  termed  allotropic  modifications.  The 
term  allotropy  is  used  both  when  the  modifications  differ  in  chemical 
as  well  as  in  physical  properties  (e.g.  oxygen  and  ozone)  and  when 
they  differ  mainly  or  entirely  as  regards  physical  properties  (e.g.  the 
modifications  of  carbon  and  of  sulphur).  The  evidence  for  the 
occurrence  of  allotropy  is  that  one  modification  can  be  converted 
into  another  without  residue,  and  that  equal  weights  of  the 
two  forms  give  identical  products  with  equal  weights  of  other 
substances.  For  example,  when  equal  weights  of  the  same  substance 
are  burned  in  oxygen  and  ozone  respectively,  equal  weights  of  the 
same  product  are  obtained.  The  different  allotropic  modifications 
of  an  element  are  characterized  by  a  difference  in  energy  content 
(pp.  135  and  241). 

The  occurrence  of  different  modifications  of  a  substance,  whether 
element  or  compound,  is  designated  by  the  general  term  isomerism^ 
and  a  distinction  is  drawn  between  physical  isomerism  or  poly- 
morphism when  .the  molecule  remains  unchanged,  and  chemical 
isomerism  when  the  molecule  is  altered.  The  chief  examples  of 
polymorphism  are  the  different  crystalline  forms  of  the  same  element 
(e.g.  sulphur,  p.  292)  or  compound  (e.g.  mercuric  iodide,  p.  475),  but 
strictly  speaking  the  different  states  of  aggregation  of  a  substance 
(solid,  liquid,  gas)  also  come  under  this  head. 

Chemical  isomerism  includes  simple  isomerism  or  metamerism  when 
two  substances  have  the  same  empirical  formula  (p.  121)  and  the 
same  molecular  weight,  but  differ  in  properties  (e.g.  ammonium 
cyanate  and  urea,  p.  343)  and  polymerism  or  polymerisation  when 
two  or  more  molecules  of  the  same  kind  combine  to  form  complex 
molecules  of  the  same  empirical  composition  (e.g.  HF  and  H2F2, 

P-  153). 

As  already  indicated,  the  term  allotropy  applies  to  all  the  isomerism 
phenomena  of  an  element,  independently  of  whether  it  is  of  a  physical 
or  of  a  chemical  character. 


CHAPTER  XIV 
OXIDES  AND  OXYGEN  ACIDS   OF  THE   HALOGENS 


eneral  —  It  has  already  been  mentioned  (p.  100)  that  although 
V_J  chlorine  does  not  combine  directly  with  free  oxygen,  three  oxides 
of  chlorine,  of  the  respective  formulae  C12O,  C1O2  and  C12O7,  can  be 
prepared  by  indirect  methods.  The  first  and  last  of  these  oxides 
are  acidic.  When  passed  into  water,  chlorine  monoxide  forms  an  acid, 
hypochlorous  acid,  HC1O,  according  to  the  equation 

C12O  +  H2O->2HC1O. 

Acids  of  this  type  are  called  oxygen  acids  or  oxyacids,  as  in  addition 
to  hydrogen  and  another  element,  in  this  case  a  halogen,  they  contain 
oxygen.  A  number  of  oxygen  acids  of  chlorine  besides  hypochlorous 
acid  are  known  ;  they  are  described  below.  From  their  property  of 
forming  acids  with  water,  acidic  oxides  are  sometimes  called  acid 
anhydrides. 

No  oxides  of  fluorine  have  been  obtained  up  to  the  present,  either 
by  direct  or  indirect  methods,  and  oxygen  acids  of  fluorine  are  also 
unknown.  No  oxides  of  bromine  are  known,  but  two  oxyacids  of 
this  element,  hypobromous  acid,  HBrO,  and  bromic  acid,  HBrO3, 
have  been  prepared.  One  stable  oxide  of  iodine,  iodine  pentoxide, 
I2O6,  and  two  oxygen  acids  are  known.  These  compounds  will  now 
be  described  in  detail. 

OXIDES  AND  OXYACIDS  OF  CHLORINE 
Three  oxides  of  chlorine  are  known,  viz.  : 

Chlorine  monoxide  (hypochlorous  anhydride;         .        C12O 
Chlorine  dioxide       .......        C1O2 

Chlorine  heptoxide  (perchloric  anhydride)      .        .       C12O7. 

Four  oxyacids  have  been  obtained  ;  the  names  and  formulae  are 
as  follows  :  — 

Hypochlorous  acid  .......  HC1O 

Chlorous  acid  .......  HC1O2 

Chloric  acid      .  ......  HC1O3 

Perchloric  acid         .......  HC1O4 

176 


OXIDES  AND   OXYGEN   ACIDS  OF  HALOGENS     177 

As  it  will  be  necessary,  in  describing  the  preparation  and  properties 
of  the  oxides,  to  mention  certain  salts  derived  from  the  oxyacids 
before  their  systematic  description  is  reached,  it  will  be  well  to  note 
here  that  the  salts  corresponding  with  hypochlorous  acid  are  known 
as  hypochlorites,  those  corresponding  with  chlorous  acid  as  chlorites, 
those  corresponding  with  chloric  acid  as  chlorates^  whilst  the  salts 
of  perchloric  acid  are  termed  perchlorates.  The  principles  on  which 
this  nomenclature  is  based  are  discussed  later. 


CHLORINE  MONOXIDE  (HYPOCHLOROUS  ANHYDRIDE),  C12O 

Preparation  —  This  compound  is  obtained  by  passing  chlorine 
over  dry  precipitated  mercuric  oxide  at  a  low  temperature  : 


The  oxide  passes  off  as  a  gas,  a  brownish  crystalline  compound, 
HgOHgCl2,  remaining  behind. 

Properties  —  Chlorine  monoxide  is  a  brownish-yellow  gas  at 
ordinary  temperatures.  It  can  readily  be  condensed  to  a  liquid, 
which  boils  at  5°.  The  oxide  is  very  unstable  ;  on  gentle  warming, 
or  in  contact  with  organic  matter,  it  breaks  up  explosively  into  its 
elements.  This  behaviour  accords  with  the  fact  that  the  monoxide 
is  a  highly  endothermic  compound  (p.  145).  On  breaking  down  into 
its  elements,  according  to  the  equation 

2C12O-X2C12+O2 

2  x  17,800  calories  are  given  out. 

The  monoxide  is  fairly  soluble  in  water,  with  which  it  combines  to 
form  hypochlorous  acid  : 

C12O  +  H2O->2HC1O. 

Composition  —  The  method  by  which  the  composition  of  the 
monoxide  was  first  established  is  instructive.  The  gas  was  collected 
in  a  tube  over  mercury  and  split  up  into  its  elements  by  gentle 
heating,  when  it  was  found  that  two  volumes  of  the  gas  gave  three 
volumes  of  the  mixture.  The  latter  was  then  treated  with  potassium 
hydroxide,  which  absorbed  the  chlorine,  leaving  one  volume  of  oxygen  ; 
the  mixture  therefore  contained  two  volumes  of  chlorine.  These 
results  can  be  satisfied  only  by  the  equation 

2C12O      ->        2C12        +         O2 
2  unit  vols.       2  unit  vols.       I  unit  vol. 
12 


iy8     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

which  indicates  that  the  gas  contains  its  own  volume  of  chlorine  and 
half  its  volume  of  oxygen. 

The  formula  C1.2O  is  confirmed  by  the  vapour-density,  which  is 
about  43.5,  corresponding  with  the  molecular  weight  87  (2  x  35.5  +  16). 


CHLORINE  DIOXIDE,  C1O2 

Preparation  —  (i)  On  gently  heating  a  mixture  of  concentrated 
sulphuric  acid  and  potassium  chlorate,  KC1O3,  chlorine  dioxide  is 
given  off  as  a  gas 


Potassium  perchlorate,  KC1O4,  and  potassium  acid  sulphate, 
KHSO4,  remain  in  the  retort.  If  the  temperature  is  allowed  to  rise 
above  a  certain  point,  the  dioxide  splits  up  into  its  elements  with  a 
violent  explosion. 

(2)  Chlorine  dioxide,  mixed  with  chlorine,  is  obtained  when 
potassium  chlorate,  KC1O3,  is  heated  with  hydrochloric  acid: 


The  mixture  of  chlorine  dioxide  and  chlorine  obtained  in  this  way 
was  regarded  by  Davy  as  a  definite  compound  of  chlorine  and  oxygen, 
and  was  called  euchlorine. 

(3)  A  steady  stream  of  chlorine  dioxide,  mixed  with  an  equal 
volume  of  carbon  dioxide,  is  obtained  by  heating  a  mixture  of 
potassium  chlorate  (40  grams)  and  oxalic  acid  (150  grams)  with  a 
little  water  (20  c.c.)  at  60°,  direct  sunlight  being  excluded  : 


Properties  —  Chlorine  dioxide  is  a  deep  yellow  gas,  with  a 
peculiar  odour,  somewhat  resembling  chlorine.  It  can  be  condensed 
to  a  deep  reddish-brown  liquid,  which  boils  at  10°,  and  on  further 
cooling  forms  yellow  crystals,  which  melt  at  —  79°.  The  dioxide  is 
extremely  unstable,  exploding  violently  on  warming  or  in  contact 
with  organic  matter  ;  it  also  decomposes  into  its  elements  on  exposure 
to  sunlight  at  the  ordinary  temperature.  Owing  to  the  readiness 
with  which  it  gives  up  oxygen,  it  is  a  powerful  oxidizing  agent.  This 
may  be  shown  very  instructively  by  adding  to  a  mixture  of  sugar  and 
potassium  chlorate  a  drop  of  concentrated  sulphuric  acid.  The 
mixture  at  once  bursts  into  flame,  the  liberated  chlorine  dioxide 


OXIDES  AND  OXYGEN   ACIDS  OF  HALOGENS     179 

igniting  the  sugar,  which  then  burns  at  the  expense  of  the  oxygen  in 
the  chlorate.     It  is  also  a  powerful  bleaching  agent. 

Chlorine  dioxide  is  readily  soluble  in  water.  When  exposed  to 
sunlight  the  dissolved  dioxide  decomposes  into  its  elements,  but  in 
the  dark  more  complicated  changes  take  place.  When  passed  into 
an  aqueous  solution  of  potassium  hydroxide,  a  mixture  of  potassium 
chlorate  and  chlorite  in  equivalent  proportions  is  formed  : 

C1O2+2KOH->KC1O2  +  KC1O3. 

The  formula  C1O2,  ascribed  to  this  compound,  is  confirmed  by 
vapour  density  determinations. 

CHLORINE  HEPTOXIDE,  C12O7 

This  compound  is  obtained  by  slowly  adding  phosphoric  pentoxide  to  per- 
chloric acid  cooled  below  - 10°,  and  then  distilling  the  mixture  on  a  water-bath. 
The  heptoxide  is  a  colourless  liquid,  which  boils  at  82°,  and  is  more  stable,  and 
therefore  a  less  powerful  oxidizing  agent,  than  the  other  oxides  of  chlorine. 


HYPOCHLOROUS  ACID,  HC1O 

Preparation — (i)  As  already  mentioned,  hypochlorous  acid  is 
obtained  by  passing  chlorine  monoxide  into  water. 

(2)  When  chlorine  is  passed  into  water  an  equilibrium  is  estab- 
lished represented  by  the  equation 

C12  +  H2O$HC1+HC1O. 

This  reaction  can  be  utilized  for  the  preparation  of  hypochlorous 
acid  by  removing  the  hydrochloric  acid,  whereby  the  equilibrium  is 
displaced  towards  the  right.  For  this  purpose  chlorine  is  passed  into 
water  in  which  a  slightly  soluble  metallic  oxide,  hydroxide,  or  carbonate 
is  suspended  (e.g.  HgO,  CaCO8,  Ag2CO3),  whereby  a  non-volatile 
chloride  is  formed  and  the  hypochlorous  acid  can  be  separated  by 
distillation. 

(3)  When  chlorine  is  passed  into  a  cold  solution  of  an  alkali,  such  as 
potassium  hydroxide,  KOH,  the  base  being  kept  in  excess  (cf.  p.  181), 
a  mixture  of  the  chloride  and  hypochlorite  of  the  base  is  obtained — 

2KOH  +  C12->KC1+KC1O  +  H2O. 

Bleaching  Powder^  which  in  solution  consists  of  a  mixture  of 
calcium  chloride  and  hypochlorite,  is  similarly  formed  by  the  action 
of  chlorine  on  moist  calcium  hydroxide,  Ca(OH)2. 


i8o     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

If  an  amount  of  dilute  nitric  acid  not  quite  equivalent  to  the  hypo- 
chlorite  is  added  to  a  mixture  of  chloride  and  hypochlorite,  the  former 
is  practically  unaffected,  whilst  hypochlorous  acid  is  set  free  from 
the  hypochlorite  according  to  the  equation 


On  distilling  the  mixture  a  dilute  solution  of  hypochlorous  acid  is 
obtained  in  the  receiver. 

Instead  of  a  mixture  of  chloride  and  hypochlorite,  a  pure  hypo- 
chlorite could  of  course  be  used  for  preparing  the  acid,  but  they 
cannot  be  obtained  so  readily. 

Properties  —  Hypochlorous  acid  has  never  been  obtained  free 
from  water.  The  dilute  aqueous  solution  is  light  yellow  in  colour,  has 
a  chlorous  smell,  and  is  moderately  stable  ;  the  concentrated  solution 
is  golden-yellow  and  is  very  unstable,  splitting  up  into  hydrochloric 
acid  and  oxygen  :  2HC1O->2HC1  +  O2. 

Owing  to  the  readiness  with  which  it  yields  oxygen,  hypochlorous 
acid  is  a  powerful  oxidizing  and  bleaching  agent.   .In  this  respect  it 
has  double  the  bleaching  capacity  of  an  equivalent  quantity  of  free 
chlorine,  as  is  evident  from  a  comparison  of  the  two  equations 
C12+H2O->2HC1  +  K>2  and  2HC1O->2HC1  +  O2. 

In  moderate  concentration  hydrochlorous  and  hydrochloric  acids 
react  to  form  free  chlorine  and  water,  represented  by  the  upper  arrow 
in  the  equation  HC1  +  HC1O$C12+H2O, 

but,  as  already  explained,  this  reaction  is  reversible,  the  position 
of  equilibrium  depending  on  the  relative  concentrations  of  the  reacting 
substances  and  other  factors.  The  removal  of  chlorine  from  the  system 
(owing  to  its  volatility  and  relatively  slight  solubility)  naturally  favours 
the  reaction  represented  by  the  upper  arrow. 

The  hypochlorites  are  derived  from  hypochlorous  acid  by  displace- 
ment of  the  hydrogen  by  metals,  and  may  be  obtained  pure  by 
neutralizing  the  free  acid  with  the  appropriate  bases.  When  heated 
alone  under  certain  conditions  they  yield  oxygen  and  the  corresponding 
chloride.1  They  are  largely  employed  for  bleaching  and  disinfecting 
purposes,  and  in  this  respect  a  solution  of  sodium  hypochlorite,  NaCIO, 
known  as  eau  de  Javelle,  and  a  mixture  of  calcium  chloride  and 
hypochlorite,  known  as  bleaching  powder,  deserve  special  mention. 

1  The  decomposition  of  hypochlorites  to  chlorine  and  oxygen  is  accelerated  by 
cobalt  oxide,  and  this  may  be  used  as  a  method  of  obtaining  oxygen. 


OXIDES  AND  OXYGEN   ACIDS  OF  HALOGENS     181 

CHLOROUS  ACID,  HC1O2 

Chlorous  acid  has  up  to  the  present  been  obtained  only  in  dilute  aqueous 
solution  by  adding  dilute  sulphuric  acid  to  a  mixture  of  potassium  chlorite 
and  chlorate  (obtained  by  passing  chlorine  peroxide  into  a  solution  of  potassium 
hydroxide,  p.  179)  and  removing  chlorine  dioxide  and  free  chlorine  by  means 
of  arsenious  acid.  A  solution  containing  free  chlorous  acid,  mixed  with 
chloric  acid,  is  thus  obtained ;  it  decomposes  fairly  rapidly  with  formation  of 
chlorine  dioxide.  Chlorous  acid  is  a  powerful  oxidizing  agent.  The  chlorites, 
on  the  other  hand,  have  only  a  weak  oxidizing  action. 

CHLORIC  ACID,  HC1O3 

Preparation — Chloric  acid  is  obtained  by  the  action  of  dilute 
sulphuric  acid  on  barium  chlorate  in  equivalent  proportions  : 
Ba(ClO3)2+H2SO4->BaSO4^  +2HC1O3. 

The  acid  is  separated  from  the  insoluble  barium  sulphate  by  filtra- 
tion and  concentrated  in  a  vacuum  over  sulphuric  acid.  In  this  way 
a  solution  containing  about  40  per  cent,  of  the  acid  is  obtained. 
When  the  attempt  is  made  to  remove  more  of  the  water  the  acid 
decomposes  into  chlorine,  oxygen,  and  perchloric  acid,  HC1O4. 

Properties — The  concentrated  acid  is  an  extremely  powerful 
oxidizing  and  bleaching  agent ;  wood  and  paper  at  once  catch  fire 
when  plunged  into  it. 

The  salts  of  chloric  acid  are  called  chlorates.  The  most  important 
salt  is  potassium  chlorate,  KC1O3,  which  is  obtained,  mixed  with 
potassium  chloride,  by  passing  chlorine  into  a  hot  concentrated 
solution  of  potassium  hydroxide,  or  by  acting  on  the  alkali  with 
ex  cess _  of  chlorine  in  the  cold  (p.  413)  : 


The  two  salts  may  be  separated  by  taking  advantage  of  the  fact 
that  the  chlorate  is  much  less  soluble  in  water  than  the  chloride. 

All  chlorates  are  soluble  in  water,  and  have  only  weak  oxidizing 
properties.  On  prolonged  heating  at  a  high  temperature  potassium 
chlorate  splits  up  completely  into  potassium  chloride  and  oxygen  : 


If,  however,  the  heating  is  stopped  at  an  intermediate  stage,  the 
mixture  is  found  to  contain  a  salt  with  more  oxygen  than  the  chlorate. 
This  salt  is  called  potassium  perchlorate,  KC1O4.  The  progress  of 


i82     A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

the  decomposition  of  the  chlorate  depends  upon  the  temperature  and 
other  factors,  but  it  is  probable  that  at  400°  the  first  stage  is  as 
follows : — 

4KC1O3->KC1  +  3KC1O4. 

At  higher  temperatures,  the  perchlorate  is  completely  decomposed 
into  chloride  and  oxygen. 

We  shall  meet  later  with  many  instances  in  which,  as  in  the 
present  case,  a  compound  of  intermediate  type  splits  up  into  one 
of  a  lower  and  another  of  a  higher  type.  The  effect  of  heating  a 
mixture  of  potassium  chloride  and  hypochlorite — 

3KC1  +  3KC10->5KC1  +  KC103, 

which  leads  to  the  same  result  as  when  chlorine  is  passed  into  hot 
potassium  hydroxide,  is  clearly  of  the  same  nature.1 

PERCHLORIC  ACID 

Preparation— (i)  A  certain  quantity  of  perchloric  acid  is  formed 
when  a  concentrated  aqueous  solution  of  chloric  acid  is  heated  (see 
chloric  acid). 

(2)  The  acid  is  usually  prepared  by  distilling  potassium  perchlorate 
with  excess  of  concentrated  sulphuric  acid  in  a  vacuum  : 

KC1O4  +  H2SO4->KHSO4  +  HClO4f. 

Properties— The  pure  acid  is  a  colourless  liquid  at  ordinary 
temperatures.  It  boils  at  16°  under  18  mm.  pressure  with  slight 
decomposition;  its  density  at  20°  is  1.7676.  With  water  it  forms 
a  constant  boiling  mixture  (p.  95)  which  contains  72.4  per  cent,  of 
acid  and  boils  at  203°  under  atmospheric  pressure.  The  acid  forms 
a  number  of  stable  compounds  with  water,  including  a  monohydrate, 
HC1O4,H2O,  and  dihydrate,  HC1O4,2H2O. 

The  concentrated  acid  is  a  powerful  oxidizing  agent,  wood  and 
paper  immediately  catching  fire  in  it.  The  dilute  acid  is  a  much 
less  powerful  oxidizing  agent  than  a  solution  of  chloric  acid  of  the 
same  concentration.  '  This  is  shown,  for  instance,  by  the  fact  that 
no  chlorine  is  given  off  when  a  dilute  solution  is  gently  warmed 
with  hydrochloric  acid.  Further,  solutions  of  perchloric  acid  have  no 
bleaching  properties. 

1  It  is,  of  course,  evident  that  only  the  hypochlorite  undergoes  change  in  this 
case:  3KC1O-»KC1O8+2KC1, 


OXIDES  AND  OXYGEN  ACIDS  OF  HALOGENS     183 

The  salts  of  this  acid,  the  perchlorates,  are  all  soluble  in  water, 
but  the  solubility  of  potassium  perchlorate  is  small,  and  advantage 
has  been  taken  of  this  to  effect  a  partial  separation  of  potassium  from 
the  other  alkali  metals. 


OXYACIDS  OF  BROMINE 

As  already  mentioned,  no  compounds  of  bromine  and  oxygen  are 
known,  but  two  oxyacids,  hypobromous  acid,  HBrO,  and  bromic 
acid,  HBrO3,  have  been  obtained.  As  their  methods  of  preparation 
and  properties  are  closely  analogous  to  those  of  the  corresponding 
chlorine  compounds,  they  can  be  dealt  with  very  briefly. 

HYPOBROMOUS  ACID,  HBrO 

Preparation—  The  acid  can  readily  be  obtained  by  shaking  up 
precipitated  mercuric  oxide  with  bromine  water  : 


On  distilling  under  reduced  pressure,  a  dilute  solution  of  the  acid 
is  obtained. 

Properties—  The  concentrated  aqueous  solution  of  hypobromous 
acid  is  straw-yellow  in  colour,  and  has  a  powerful  bleaching  and 
oxidizing  action.  On  heating,  it  readily  splits  up  into  hydrobromic 
acid  and  oxygen.  The  hypobromites  are  also  oxidizing  agents. 
Sodium  hypobromite,  which  is  the  best-known  of  these  salts,  forms 
a  yellow  solution  with  water,  and  is  employed  as  a  mild  oxidizing 
agent  in  organic  chemistry. 

BROMIC  ACID,  HBrO3 

Preparation—  (i)  By  treating  barium  bromate,  Ba(BrO3)2,  with 
the  calculated  quantity  of  sulphuric  acid  : 


(2)  By  passing  excess  of  chlorine  through  bromine  water  : 
Br2  +  5Cl2+6H2O->2HBrO3+ioHCl. 

In  this  interesting  reaction,  bromine,  though  itself  an  oxidizing 
agent  under  certain  conditions,  is  oxidized  by  chlorine,  a  more 
powerful  oxidizing  agent. 


1  84     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

(3)  By  treating  the  sparingly  soluble  silver  bromate,  AgBrO3,  with 
bromine  : 

5  AgBrO3  +  3Br2  -f  sH2O->5  AgBrj  +  6HBrO3. 


Properties  —  Bromic  acid  is  known  only  in  aqueous  solution. 
On  heating  it  splits  up  into  bromine,  water  and  oxygen.  It  has 
oxidizing  and  bleaching  properties. 

The  corresponding  salts,  the  bromates,  are  mostly  soluble  with 
difficulty  in  water.  Those  of  the  alkalis  decompose  into  bromide 
and  oxygen  on  heating,  without  the  intermediate  formation  of  per- 
bromate. 

OXIDE  AND  OXYACIDS  OF  IODINE 

Hypoiodous  Acid  —  When  iodine  is  added  to  a  cold  solution  of  potassium 
hydroxide  a  colourless  solution  is  obtained.  When  freshly  prepared  the  solu- 
tion bleaches  indigo,  and  on  addition  of  weak  acids,  such  as  acetic  acid,  iodine 
is  set  free.  If,  however,  the  solution  is  kept  for  a  considerable  time,  or  is  heated 
and  allowed  to  cool,  it  no  longer  shows  the  above  properties.  The  simplest 
explanation  of  these  observations  is  that  the  reaction  between  iodine  and  cold 
potassium  hydroxide  is  similar  to  that  with  chlorine  : 

I2+2KOH-»KI  +  KIO  +  H2O, 

and  that  the  hypoiodite,  KIO,  has  a  bleaching  action.  Weak  acids  set  free 
hydriodic  and  hypoiodous  acid,  which  react  immediately  with  liberation  of 
iodine  : 

HI  +  HIO->I2+H2O. 

The  mixture  of  alkali  iodide  and  hypoiodite  is,  however,  unstable,  and  a 
mixture  of  iodide  and  iodate  results  on  keeping  or  on  heating  the  solution  :  J 

3KI  +  3KIO-»sKI  +  KI03. 

The  iodate,  KIO3,  yields  oxygen  much  less  readily  than  the  hypoiodite,  and 
the  solution  containing  it  has  therefore  no  bleaching  properties. 

IODINE  DIOXIDE,  IO2    , 

The  existence  of  a  compound  of  this  formula  was  first  mentioned  by  Millon,  and 
quite  recently  Pattison   Muir   has  described  its  preparation  by  heating  iodic    ' 
acid  with  sulphuric  acid.     It  occurs  in  small  lemon-yellow  crystals,  which  begin 
to  decompose  at  130°,  and  on  further  heating  decompose  completely  into  iodine 
and  oxygen. 

IODINE  PENTOXIDE,  I2O6 

Preparation  —  Iodine  pentoxide  is  the  anhydride  of  iodic  acid, 
and  is  obtained  by  heating  this  acid  at  199  to  200°  : 

2HIO3->I2O6+H2O. 


OXIDES  AND   OXYGEN   ACIDS  OF  HALOGENS     185 

Properties  —  It  is  a  white  crystalline  substance,  which  splits  up 
completely  into  iodine  and  oxygen  on  heating  at  300°,  and  dissolves 
in  water  to  form  iodic  acid. 

IODIC  ACID,  HIO3 

Preparation—  (i)  By  treating  barium  iodate  with  the  calculated 
amount  of  dilute  sulphuric  acid,  filtering,  and  evaporating  in  vacuo  : 

Ba(IO3)2+H2SO4->BaSO4|  +  2HIOs. 

(2)  Iodic  acid  is  readily  obtained  by  the  oxidation  of  iodine  in 
presence  of  water,  for  example,  by  boiling  iodine  with  nitric  acid  till 
the  red  colour  disappears  : 

3l2+  ioHNO3->6HIO3  +  ioNOf  +  2H2O. 

Instead  of  nitric  acid,  chlorine  may  be  used  as  oxidizing  agent 
(cf.  p.  90). 

Properties  —  Iodic  acid  occurs  in  colourless  crystals,  which 
dissolve  in  water  to  form  a  strongly  acid  solution.  Iodic  acid  is  an 
oxidizing  agent,  but  not  so  energetic  as  chloric  or  bromic  acids  ; 
when  brought  in  contact  with  hydriodic  acid,  HI,  the  hydrogen  of 
the  latter  is  oxidized  to  water,  and  iodine  and  water  are  the  sole 
products  : 


When  iodic  acid  is  heated  to  IK>°,  it  loses  part  of  its  water,  and 
yields  the  compound  HI3O8: 

3HI03-H20->HI308. 

At  190°  to  200°  all  the  water  is  driven  off  and  iodine  pentoxide 
remains.  The  iodates,  like  the  chlorates  and  bromates,  are  obtained 
by  the  action  of  the  free  halogen  on  a  hot  solution  of  the  base,  and 
can  be  separated  from  the  iodides,  which  are  simultaneously  formed, 
by  taking  advantage  of  the  smaller  solubility  of  the  iodates. 

PERIODIC  ACID,  HIO4,2H2O  or  H6IO6 

Preparation  —  (i)  This  compound  is  obtained  by  the  action  of 
sulphuric  acid  on  barium  periodate  : 

Ba(IO4)2  +  H2SO4+2  H2O->BaSO4  +  2  HIO4(H2O)2. 
(2)  Alkali  periodates  are  obtained  by  the  electrolytic  oxidation  of 


i86     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

alkali  iodates.  By  the  same  method  periodic  acid  can  be  obtained 
from  iodic  acid. 

(3)  A  periodate  of  sodium  is  obtained  by  the  action  of  chlorine  on 
sodium  iodate  in  hot  alkaline  solution. 

Properties — The  acid  is  obtained,  on  evaporation  of  its  aqueous 
solution,  as  colourless  crystals  of  the  composition  HIO4,2H2O,  which 
melt  at  130°,  and  on  heating  at  a  slightly  higher  temperature  split  up 
into  iodine  pentoxide,  water,  and  oxygen.  The  anhydrous  acid, 
HIO4,  has  not  been  obtained,  but  salts,  such  as  AgIO4  and  NaIO4, 
derived  from  it,  are  known  (see  next  section). 

Anhydrides,  Acids,  and  Hydrates  of  Acids— An  acid  has  been  provisionally 
defined  as  a  substance  containing  hydrogen,  the  whole  or  part  of  which  can  be 
replaced  by  metals,  with  formation  of  salts.  An  anhydride  or  acidic  oxide  is  a 
compound  derivable  from  an  oxyacid  by  abstraction  of  one  or  more  molecules  of 
water,  and,  conversely,  the  formula  of  an  oxyacid  is  derived  from  that  of  the 
anhydride  by  addition  of  one  or  more  molecules  of  water. 

When  the  anhydride  corresponding  with  an  acid  is  known,  the  latter  cannot 
always  be  obtained  directly  from  the  former,  or  the  former  from  the  latter.  In 
the  halogen  group,  as  we  have  seen,  these  changes  can  generally  be  accomplished, 
but  we  shall  see  later  that  the  acid  corresponding  with  nitrous  oxide,  N2O,  cannot 
be  obtained  directly  from  the  latter,  nor  can  phosphorus  pentoxide,  P2O5,  be 
obtained  directly  from  the  corresponding  acid. 

In  many  cases  the  anhydrides  corresponding  with  oxyacids  have  not  so  far 
been  obtained.  Thus,  from  analogy,  the  existence  of  an  anhydride,  I2O7,  corre- 
sponding with  periodic  acid,  HIO4,  would  be  anticipated,  but  so  far  it  is  unknown. 
Many  other  examples  of  this  have  been  met  with  in  the  present  chapter.  We 
shall,  however,  sometimes  find  it  convenient,  for  purposes  of  formulation,  to 
assume  the  existence  of  these  hypothetical  oxides. 

As  acids  are  derived  from  acidic  oxides  by  addition  of  water,  it  is  conceivable 
that  one  oxide,  by  association  with  different  amounts  of  water,  may  give  rise  to 
more  than  one  oxyacid.  That  this  is  the  case  we  have  already  learnt  in  con- 
nexion with  iodic  acid  (p.  185).  For  purposes  of  comparison,  the  acid  itself  may 
be  written,  I2O5,H2O  (  =  2HIO3),  and  the  other  acid,  obtained  by  heating  the 
ordinary  acid  to  110°,  as  3l2O5,H2O  (=2  HI3O8). 

It  has  been  found  that  salts  of  different  types  are  obtained  from  solutions  con- 
taining periodates,  for  example,  AgIO4  and  Ag4I2O9,  and  these  may  be  regarded 
as  being  derived  from  acids  of  the  types  HIO4,  H5IO6,  H4I2O9,  etc.  It  is  not  at 
first  sight  evident  that  these  acids  are  all  periodic  acids,  but  further  consideration 
shows  that  they  may  be  regarded  as  being  formed  by  the  association  of  different 
quantities  of  water  with  the  hypothetical  periodic  anhydride,  I2O7,  thus,  I2O7,H2O 
=2HIO4;  I2O7,5H2O=2H5IO6,  and  I2O7,2H2O=H4I2O9.  It  is  interesting  to 
note  that,  as  already  pointed  out,  the  stable  acid  is  not  HIO4,  but  H5IO6  (which 
may  be  written  HIO4,2H2O),  and  the  majority  of  the  salts  of  periodic  acid  are 
derived  from  the  latter  acid.  On  the  other  hand,  the  salts  derived  from  perchloric 
acid,  HC1O4,  are  formed  by  the  displacement  of  one  hydrogen  only  by  metals, 
although  a  number  of  compounds  of  the  acid  and  water  are  known  (p.  182).  We 


OXIDES  AND  OXYGEN  ACIDS  OF  HALOGENS     187 

must  therefore  suppose  that  the  water  in  the  dihydrate  of  perchloric  acid, 
HC1O4,2H2O,  is  less  intimately  associated  with  the  rest  of  the  molecule  than  in 
the  case  of  the  dihydrate  of  periodic  acid,  as  in  the  former  case  none  of  the  hydro- 
gens is  displaceable  by  metals.  For  this  reason  the  water  associated  with  per- 
chloric acid  may  appropriately  be  regarded  as  water  of  crystallization ;  in  the 
case  of  periodic  acid  as  water  of  constitution. 

As  already  mentioned,  acids  of  which  only  one  hydrogen  is  displaceable  by 
metals  are  termed  monobasic,  and  nearly  all  the  acids  hitherto  met  with  are  of 
this  type.  Acids  which  contain  more  than  one  displaceable  hydrogen  are  said  to 
be  polybasic,  and  some  of  the  complex  periodic  acids  just  considered  are  of  this 
type.  This  question  will  be  fully  Considered  later  in  connexion  with  the  oxyacids 
of  phosphorus. 

Strong  and  Weak  Acids— We  have  seen  that  hypochlorous 
acid  can  be  obtained  by  adding  nitric  acid  to  a  solution  of  a  hypo- 
chlorite,  for  example,  potassium  hypochlorite,  and  distilling  : 

KC1O  +  HNO3-»KNO3+  HClOf. 

One  reason  why  the  reaction  proceeds  practically  to  completion  in 
the  direction  of  the  arrow  is  that  hypochlorous  acid  is  volatile  and 
escapes  from  the  system,  but  it  is  of  interest  to  inquire  what  occurs 
when  equivalent  proportions  of  potassium  hypochlorite  and  nitric 
acid  are  mixed  in  the  cold,  under  which  condition  practically  all  the 
hypochlorous  acid  remains  in  the  system.  As  would  be  anticipated,  an 
equilibrium  is  then  established,  represented  by  the  equation 

KClO  +  HNO3^tKNO3+HClO, 

and  it  can  be  shown,  by  methods  which  it  would  lead  too  far  to 
describe,  that  it  lies  very  near  the  right-hand  side ;  in  other  words, 
nitric  acid  displaces  hypochlorous  acid  almost  completely  from  com- 
bination with  a  base.  In  such  a  case  we  are  accustomed  to  state  that 
nitric  acid  is  a  much  stronger  acid  than  hypochlorous  acid.  The  above 
is  one  of  the  standard  methods  by  which  the  relative  strengths  of  two 
acids  is  compared.  They  are  allowed  to  compete,  under  such  condi- 
tions that  all  the  reacting  substances  remain  in  the  system,  for  an 
amount  of  base  insufficient  to  saturate  both  of  them,  and  the  ratio  in 
which  the  distribution  takes  place  is  measured.  The  stronger  acid  is 
the  one  which  takes  possession  of  most  of  the  base. 

Of  the  acids  so  far  considered,  hydrochloric,  hydrobromic  and 
hydriodic  acids  are  among  the  strongest,  whilst  hydrofluoric  acid 
is  considerably  weaker.  Most  of  the  oxyacids  of  the  halogens  are 
strong  acids,  but,  as  already  mentioned,  hypochlorous  acid  is  an 
exception. 


1 88     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

In  connexion  with  the  relative  strengths  of  acids  the  effect  of 
differences  of  volatility  and  of  solubility  in  disturbing  the  equilibrium 
must  be  borne  in  mind.  Thus  sulphuric  acid  is  actually  a  weaker 
acid  than  hydrochloric  acid,  but  on  account  of  its  volatility  the  latter 
is  displaced  almost  completely  from  combination  by  the  former  on 
heating. 

Nomenclature  —  The  system  adopted  in  naming  the  halogen 
oxyacids  is  in  general  use  for  all  oxyacids,  and  will  now  be  briefly 
described.  It  has  already  been  explained  that  the  names  of  binary 
compounds — those  containing  two  elements  only — end  in  ide. 

The  name  of  the  most  important  oxyacid  usually  ends  in  ic — 
example,  chloric  acid,  phosphoric  acid.  Acids  containing  more 
oxygen  have  the  prefix  per,  e.g.  perchloric  acid,  whilst  the  names 
of  acids  containing  less  oxygen  end  in  ous,  e.g.  chlorous  acid,  sul- 
phurous acid.  Those  containing  still  less  oxygen  than  the  ous  acid 
have,  in  addition,  the  prefix  hypo,  e.g.  hypochlorous  acid. 

The  names  of  the  salts  derived  from  the  ic  acids  ends  in  ate,  e.g. 
potassium  chlorate,  potassium  sulphate,  those  derived  from  the  ous 
acids  end  in  ite,  e.g.  potassium  chlorite,  potassium  sulphite.  The 
salts  of  the  per  .  .  .  ic  acids  are,  of  course,  called  per  .  .  .  ates,  e.g. 
potassium  perchlorate,  and  those  of  the  hypo  .  .  .  ous  acid  are  called 
hypo  .  .  .  ites,  e.g.  potassium  hypochlorite. 

The  same  principles  are  used  in  naming  the  oxides.  If  there  is 
only  one,  the  name  of  the  element  followed  by  oxide  is  generally 
used,  e.g.  magnesium  oxide.  If  there  are  two,  the  name  of  that  con- 
taining less  oxygen  ends  in  ous,  that  containing  the  higher  proportion 
of  oxygen  in  ic,  e.g.  mercurous  oxide,  mercuric  oxide.  An  oxide  with 
more  oxygen  than  the  ic  oxide  has  the  prefix  per,  e.g.  barium 
peroxide,  one  with  less  oxygen  than  the  ous  oxide  has  the  prefix 
hypo,  e.g.  hypochlorous  oxide.  Some  exceptions  to  these  rules  will 
be  met  with  in  the  course  of  our  work. 

Another,  and  in  some  respects  preferable,  system  is  to  indicate  the 
number  of  atoms  of  oxygen  by  prefixing  the  corresponding  Latin  or 
Greek  number,  for  example,  chlorine  monoxide,  chlorine  dioxide, 
iodine  pentoxide. 

Valency  in  the  Halogen  Group— The  conception  of  valency 
has  already  been  explained,  and  we  have  seen  that  it  is  not  usually 
constant  for  any  one  element,  but  depends  on  the  nature  of  the  other 
elements  with  which  it  is  combined.  The  variable  nature  of  valency 
is  well  illustrated  by  the  halogen  compounds  discussed  in  this 
chapter. 


OXIDES  AND  OXYGEN  ACIDS  OF  HALOGENS     189 

With  regard  to  hydrogen,  chlorine,  bromine  and  iodine  are  invari- 
ably univalent,  the  only  known  compounds  having  the  graphic  formulae 
H  — Cl,  H  — Br,  H  — I.  Hydrogen  fluoride,  however,  appears  to  have 
the  formula  H2F2  under  certain  conditions.  In  this  case  fluorine  may 
perhaps  be  trivalent,  H  —  F  =  FH.  The  existence  of  compounds  such  as 
IC13,  IF6,  and  BrF3  definitely  proves  the  multivalency  of  the  halogens. 

When  the  attempt  is  made  to  measure  valencies  by  means  of  oxygen 
compounds,  uncertainty  arises  from  the  possibility  of  direct  linkings 
between  two  atoms  of  oxygen.  The  graphic  formula  of  chlorine 

Cl\ 

monoxide  may  be  written  thus       yO  and  that  of  the  corresponding 

acid   as         >O,    the    chlorine    being    umvalent    and    the    oxygen 

CK 

divalent.  When,  however,  we  come  to  chlorine  dioxide,  C1O2,  diffi- 
culties arise.  The  chlorine  could  be  represented  as  divalent, 

/O  ,0 

Cl<    I  ,  or  as   quadrivalent,    Cl£     ,  but   as   we  have  no  means  of 

X0  X0 

deciding  the  question,  there  is  nothing  gained  by  discussing  the 
matter  further. 

The  oxyacids  of  chlorine  have  the  formulae  HC1O,  HC1O2,  HC1O3, 
and  HC1O4  respectively.  These  can  be  represented  graphically  as 
follows:- 

Cl— OH;  O  =  C1— OH;      >C1— OH  ;       J:C1<        , 
O^  O^     XOH 

the  valency  of  the  chlorine  being  one,  three,  five,  and  seven  respec- 
tively. The  corresponding  compounds  of  bromine  and  iodine  can  be 
represented  in  a  similar  way. 

It  must  be  admitted  that  the  evidence  in  favour  of  the  above  graphic  formulas 
for  the  oxyacids  is  rather  slender.  On  the  usual  assumption  that  the  valency  of 
an  element  is  the  same  in  an  ox^de  and  the  corresponding  oxyacid,  the  formulae 
of  chlorous  oxide,  chlorine  pentoxide  (both  hypothetical),  and  chlorine  heptoxide 

O<v  *O  O^  ^O 

would  be  O=C1—  O— C1=O,       J^Cl— O— Clf     ,  and  O^Cl— O— C1==O    re- 

O^  \Q  o/  \O 

spectively.  The  fact  that  potassium  perchlorate  and  permanganate  are  isomorphous 
and  that  both  elements  belong  to  the  seventh  group  of  the  periodic  table  (p.  380) 
lends  some  support  to  the  view  that  chlorine  can  act  as  a  septivalent  element. 
Another  possibility  is  that  the  halogens  are  after  all  univalent,  and  that  the  oxygen 
atoms  are  arranged  in  chains  thus,  Cl— O— O— O— Cl,  Cl— O— O— O— O— O— Cl 
and  Cl— O— O— O— O— O— O— O— Cl  with  analogous  formulae  for  the  cor- 
responding acids,  thus  Cl— O — O — O— OH  would  represent  percnloric  acid.  An 
objection  to  these  formulae  is  that  chains  of  atoms  of  the  same  kind  are  not  usually 
stable,  whilst  perchloric  acid  is  quite  a  stable  compound.  A  suggestion  which 
has  been  made  is  that  the  oxygen  atoms  are  arranged  to  form  a  ring  thus  :— 

/O 
H-O— C1-O<   | 

N) 


190     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Compounds  of  the  Halogens  with  each  other—  Iodine  forms  compounds 
with  each  of  the  other  halogens,  and  a  compound  of  chlorine  and  bromine  is 
also  known. 

Iodine  monochloride,  IC1,  is  obtained  in  the  form  of  a  dark  reddish-brown  liquid 
by  passing  dry  chlorine  over  iodine.  It  has  been  obtained  in  two  crystalline 
forms,  which  melt  at  27.  2°  and  13.9°  respectively.  It  is  decomposed  by  water, 
with  formation  of  hydrochloric  and  iodic  acids  and  liberation  of  iodine  : 


Iodine  trichloride,  IC13,  is  obtained  by  the  further  action  of  chlorine  on  iodine 
monochloride  or  by  the  action  of  chlorine  in  excess  on  iodine.  It  occurs  in 
orange  needles,  and  dissociates  on  heating: 


Iodine  trichloride  is  soluble  in  water,  and  the  solution  is  used  as  an  antiseptic. 

Iodine  monobromide  ,  IBr,  is  a  crystalline  compound  which  melts  at  36°  and  is 
decomposed  by  water  with  liberation  of  iodine. 

Iodine  pentafluoride,  IF5,  is  obtained  by  direct  combination  of  the  elements 
(Moissan,  1902),  and  is  a  colourless  liquid  which  boils  at  97°  and  decomposes  on 
heating  to  400°. 

Bromine  trifluoride,  BrF8,  obtained  by  direct  combination  of  its  elements, 
(Prideaux,  Lebeau,  1905),  is  a  colourless  liquid  which  reacts  violently  with  water, 
liberating  oxygen  and  forming  chiefly  hypobromous  and  hydrofluoric  acids. 

Comparison  of  the  Halogens  and    Summary  —  As  the 

halogens  are  the  first  family  of  elements  we  have  considered  in 
detail,  a  comparison  of  their  behaviour  is  of  great  interest.  The 
most  important  fact  brought  out  by  this  comparison  is  that  there 
is  considerable  resemblance  between  the  properties  of  the  elements 
themselves,  and  between  those  of  their  corresponding  compounds, 
and,  further,  there  is  a  gradual  variation  both  of  physical  and 
chemical  properties  corresponding  with  increasing  atomic  weight. 
Thus,  just  as  bromine  is  intermediate  Jo  chlorine  and  iodine  as 
regards  atomic  weight,  it  is  also  intermediate  as  regards  physical 
and  chemical  properties.  This  generalization  is  illustrated,  as 
regards  physical  properties,  in  the  accompanying  table. 


Property. 

Fluorine. 

Chlorine. 

Bromine. 

Iodine. 

Atomic  weight    . 
Melting-point 
Boiling-point  .     . 
Density      .     .     . 

Colour  .... 

19 
-223° 
-187° 

/pale  greenish 
I         yellow 

3S-46 

-102° 

-34° 
I.3S 

greenish 
yellow 

79.96 
-7° 
59° 
3-*9 
reddish 
brown 

126.85 

+  H4* 

184° 
4.95  (solid) 

bluish  black 

OXIDES  AND  OXYGEN  ACIDS  OF  HALOGENS     191 

Without  exception,  the  magnitude  of  the  physical  properties  increase, 
or  become  more  pronounced,  as  the  atomic  weight  increases. 

The  same  is  true  of  the  chemical  behaviour.  Thus  the  affinity  for 
hydrogen  is  greatest  in  the  case  of  fluorine,  and  gradually  diminishes 
until,  in  the  case  of  iodine,  the  rate  of  combination  is  very  slow  and 
is  incomplete.  This  is  perhaps  best  shown  by  a  comparison  of  the 
heats  of  formation  of  the  hydrogen  compounds  (in  gaseous  form) 
from  their  elements  (p.  145). 

H2+F2  =2HF  +  2  x  38,600  cal. 

H2+C12=2HC1  +  2X  22,000  „ 

H2  +  Br2=2HBr  +  2x8,400    „ 
=2HI-2X6,ioo      „ 


The  order  of  affinity  of  the  halogens  for  metals  is  the  same  as  for 
hydrogen,  as  is  shown  by  the  fact  that  the  element  of  smaller  atomic 
weight  displaces  that  with  a  larger  atomic  weight  from  combination. 

In  contrast  to  the  behaviour  with  hydrogen  and  the  metals,  the 
affinity  for  oxygen  increases  with  the  atomic  weight  of  the  halogen. 
Thus  fluorine  forms  no  compound  with  oxygen,  the  compounds  of 
chlorine  and  oxygen  are  unstable,  whilst  iodine  pentoxide  is  rela- 
tively stable.  Similarly,  the  halogens  of  higher  atomic  weight,  under 
certain  circumstances,  displace  the  others  from  combination  in  oxygen 
compounds,  thus  potassium  chlorate  is  transformed  almost  com- 
pletely into  potassium  bromate  by  fusing  with  potassium  bromide  : 

KC1O3+  KBr->KBrO3+KCl. 

The  more  important  points  of  resemblance  in  the  halogen  group 
is  that  all  form  compounds  of  the  type  HX  with  hydrogen  (X  =  halo- 
gen) ;  these  compounds  are  colourless  gases,  which  fume  in  the  air, 
and  form  strong  acids  when  dissolved  in  water.  Further,  the  salts 
containing  the  same  metals  and  different  halogens  are  of  the  same  type, 
and  the  compounds  with  the  alkali  metals  occur  in  cubic  crystals. 
It  is  of  interest  to  note,  however,  that  the  behaviour  of  fluorine 
differs  more  from  that  of  the  other  halogens  than  the  latter  do 
among  themselves.  Thus  hydrogen  fluoride  tends  to  polymerize 
(p.  153),  double  salts  of  the  type  KF'HF  exist,  and  silver  fluoride, 
unlike  the  other  silver  halides,  is  readily  soluble  in  water. 


CHAPTER  XV 

OSMOTIC   PRESSURE   AND    MOLECULAR   WEIGHT 
IN   SOLUTION 


eneral  —  When  a  compound  can  be  converted  into  vapour  with- 
out  decomposition,  its  molecular  weight  can  be  determined  by 
the  volumetric  method  already  described  (p.  115).  Many  compounds, 
however,  cannot  be  vaporized  without  decomposition,  and  their 
molecular  weights  cannot  therefore  be  determined  by  the  standard 
method.  An  example  of  this  already  met  with  is  hydrogen  peroxide. 
Analysis  shows  that  it  contains  hydrogen  and  oxygen  in  equal  atomic 
proportions,  so  that  its  molecular  formula  is  either  HO  or  an  integral 
multiple  of  HO,  such  as  H2O2,  H3O3,  etc.  A  more  or  less  definite 
decision  between  these  formulae  can  be  arrived  at  on  the  ground  of 
its  chemical  behaviour.  For  example,  the  readiness  with  which  it 
splits  up  into  water  and  oxygen  appears  to  be  most  in  harmony  with 
the  generally  accepted  formula.  It  will,  however,  clearly  be  of  great 
advantage  in  this  and  similar  cases  if  the  molecular  weight  can  be 
determined,  in  solution,  for  instance,  by  a  direct  method.  Within  the 
last  thirty  years  such  methods  have  been  devised,  mainly  owing  to 
the  labours  of  the  Dutch  chemist,  van  't  Hoff,  and  have  contributed 
enormously  to  the  advance  of  chemistry.  These  methods  depend 
ultimately  on  the  conception  of  osmotic  pressure,  which  will  now  be 
^considered. 

Osmosis.  Osmotic  Pressure—  A  piece  of  parchment  paper  or 
animal  membrane  is  tied  over  the  mouth  of  a  thistle  funnel  A,  which 
is  then  inverted,  filled  up  to  a  point  on  the  stem  with  an  aqueous 
solution  of  copper  sulphate,  and  then  supported  in  a  vessel  of  water 
B,  so  that  the  liquid  outside  and  inside  is  at  the  same  height  (Fig.  45). 
After  a  time  it  will  be  observed  that  the  solution  has  risen  to  a  higher 
point  in  the  tube  and  the  level  of  the  water  has  slightly  fallen. 
Further,  the  water  in  the  vessel  is  slightly  blue,  showing  that  a  little 
copper  sulphate  has  passed  through  the  membrane. 

The  phenomenon  is  clearly  analogous  to  the  interdiffusion  of  two 
gases  separated  by  a  porous  membrane  (p.  46).  Water  passes  inte 

192 


OSMOTIC    PRESSURE 


193 


the  tube  more  quickly  than  the  copper  sulphate  diffuses  out,  with  the 
result  that  the  bulk  of  solution  inside  increases.  The  process  is 
called  osmosis.  The  inward  movement  is  sometimes 
termed  endosmosis^  and  the  outward  stream  exosmosis. 

When  a  bladder  is  used  as  membrane  one  of  the 
substances  (water)  diffuses  through  it  much  faster  than 
the  other  (copper  sulphate),  and  it  might  be  suggested 
that  with  a  suitable  membrane  the  passage  outwards 
of  the  dissolved  substance  might  be  stopped  altogether 
without  interfering  with  the  entry  of  water.  This  has 
been  actually  realized  with  a  membrane  of  copper 
ferrocyanide,  when  cane  sugar  is  used  as  solute.  A 
membrane  which  allows  one  substance  to  pass  through 
and  entirely  prevents  the  passage  of  another  is  said 
to  be  semipermeable. 

The  copper  ferrocyanide  membrane  is  usually  pre-  FlG'  45< 
pared  by  interaction  of  solutions  of  copper  sulphate  and  of  potassium 
ferrocyanide.  It  is  much  too  weak  to  use  in 
place  of  the  bladder  in  the  experiment  described 
above  ;  but  this  difficulty  can  be  overcome  by 
depositing  it  in  the  walls  of  a  porous  pot,  as 
first  suggested  by  the  German  botanist  Pfeffer 
(1877).  The  pot  (one  of  those  commonly  used 
for  experiments  in  gas  diffusion  is  suitable)  is 
first  carefully  washed,  soaked  in  water  for  some 
time,  then  nearly  filled  with  a  solution  of  copper 
sulphate  (2.5  grams  per  litre)  dipped  nearly  to 
the  neck  in  a  solution  of  potassium  ferrocyanide 
(2.1  grams  per  litre),  and  allowed  to  stand  for 
some  hours.  The  salts  diffuse  through  the 
walls  of  the  pot,  and  at  their  junction  form  a 
membrane  of  copper  ferrocyanide  which,  owing 
to  its  being  supported  by  the  walls  of  the  vessel, 
is  capable  of  withstanding  fairly  high  pressures. 
The  cell  is  then  taken  out,  washed,  rilled  with 
a  concentrated  solution  of  sugar,  and  closed 
with  a  well-fitting  cork  through  which  passes  a 
glass  tube  open  at  both  ends  (Fig.  46).  When 
FIG.  46.  a  cen  tnus  prepared  is  immersed  in  water  the 

latter  passes  in  and  the  lave)  of  liquid  in  the  narrow  tube  slowly  rises, 
and  finally  reaches  a  point  at  which,  if  the  cell  has  been  properly 
13 


B 


i94     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

prepared,  it  remains  constant  for  days.  Under  these  circumstances 
the  pressure  inside  the  cell  is  greater  than  that  outside  by  an  amount 
measured  by  the  difference  in  height  of  the  liquid  inside  and  outside 
the  cell.  This  excess  of  pressure,  which  must  prevail  inside  the  cell  in 
order  to  prevent  more  water  flowing  in  through  the  semi-permeable 
membrane,  is  termed  the  osmotic  pressure  of  the  solution. 

If  the  cell  is  dipped  into  a  more  concentrated  solution  of  sugar 
instead  of  into  water,  water  passes  outwards  from  the  more  dilute  to 
the  more  concentrated  solution,  so  that  the  concentrations  tend  to 
become  equal.  In  the  same  way,  if  solutions  of  different  substances 
are  separated  by  a  semi-permeable  membrane,  the  solvent  always 
passes  through  the  membrane  from  the  solution  with  lower  to  that 
with  higher  osmotic  pressure.  In  other  words,  the  direction  of  flow 
is  always  such  that  the  pressures  on  the  two  sides  of  the  membrane 
tend  to  become  equal. 

The  exact  mode  in  which  the  osmotic  pressure  arises  is  not  well 
understood.1  According  to  one  view,  the  extra  pressure  which  pre- 
vails inside  the  cell  is  due  to  the  bombardment  of  the  interior  walls 
by  the  particles  of  solute  ;  just  as,  according  to  the  kinetic  theory 
(p.  48),  gas  pressure  is  due  to  impacts  of  the  particles  on  the  walls  of 
the  containing  vessel.  It  is  interesting  to  note  that,  just  as  a  mem- 
brane of  copper  ferrocyanide  allows  water,  but  riot  dissolved  sugar, 
to  pass  through,  a  membrane  made  of  metallic  palladium  allows 
hydrogen  to  pass  through,  but  is  impervious  to  other  gases.  On  this 
principle  an  apparatus  can  be  constructed  which  admits  of  the  measure- 
ment of  the  partial  pressure  of  a  gas  such  as  nitrogen,  when  mixed 
with  hydrogen.  A  vessel  of  palladium,  which  can  be  heated  to  any  con- 
venient temperature,  contains  nitrogen  at  a  pressure  of  say  half  an 
atmosphere,  as  observed  on  a  manometer.  When  the  vessel,  heated 
to  600°,  is  surrounded  by  hydrogen  at  atmospheric  pressure,  the  latter 
passes  inwards  through  the  membrane  and  the  pressure  inside  in- 
creases. When  equilibrium  is  finally  attained,  the  pressure  inside  the 
cell  is  about  \\  atmospheres.  The  excess  of  pressure  inside  over  that 
outside  is  clearly  due  to  the  nitrogen.  It  is  evident  that  there  is  a 
close  analogy  between  this  experiment  and  that  by  means  of  which 
the  osmotic  pressure  of  dissolved  cane  sugar  is  determined. 

If  a  concentrated  solution  of  copper  sulphate  is  carefully  placed, 

by  means  of  a  pipette,  at  the  bottom  of  a  beaker  of  water  so  that  there 

is  a  definite  boundary  between  the  two  layers,  it  will  be  found  after  a 

time  that  the  copper  sulphate  is  uniformly  distributed  through  the 

1  Cf.  Physical  Chemistry,  p.  106. 


OSMOTIC   PRESSURE  195 

water  (diffusion  of  liquids).  We  may  in  this  case  assume  that  the 
cause  of  diffusion  is  the  higher  osmotic  pressure  of  the  copper  sulphate 
in  certain  parts  of  the  system  (just  as  the  cause  of  gaseous  diffusion 
is  the  unequal  partial  pressures  of  the  gases  in  different  parts  of  the 
system),  and  equilibrium  is  attained  when  the  concentration  (and  the 
osmotic  pressure)  of  the  solute  is  uniform  throughout. 

Osmotic  Pressure  and  Molecular  Weight  in  Solution 
— Regarding  osmotic  pressure  as  analogous  to  gas  pressure,  the 
relationship  between  volume,  osmotic  pressure  and  temperature  can 
now  be  investigated  for  solutions,  as  has  already  been  done  for 
gases. 

Firstly,  it  may  be  shown  experimentally  that  the  osmotic  pressure  is 
approximately  proportional  to  the  concentration  of  the  solution.  This 
is  evident  from  the  following  results  obtained  by  Pfeffer  for  solutions 
of  cane  sugar,  in  which  the  concentrations  are  expressed  in  grams  per 
100  c.c.  of  solution  and  the  pressures  in  cm.  of  mercury. 

Concentration,  C  .         .        .        .         I          2          4  6 

Osmotic  pressure,  P               .        -53-1     101.6     208.2     307.5 
Ratio,  P/C 53.5        50.8       52.1       51.3 

As  the  concentration  is  inversely  proportional  to  the  Volume  V  in 
which  a  definite  quantity  of  solute  is  dissolved,  we  obtain,  by  substi- 
tuting i/V  for  C,  the  relationship  PV  =  constant,  that  is,  the  osmotic 
pressure  exerted  by  a  definite  quantity  of  a  solute  x  the  volume  of 
solution  in  which  it  is  dissolved  is  constant  at  constant  temperature 
— a  result  exactly  analogous  to  Boyle's  law  for  gases. 

Further,  it  may  be  shown  that  the  osmotic  pressure,  like  the  gas 
pressure,  increases  proportionately  to  the  absolute  temperature  T  at 
constant  volume.  Pfeffer  found  the  osmotic  pressure  of  a  I  per  cent, 
solution  of  cane  sugar  to  be  51.0  cm.  at  14.2°.  If  P  is  proportional  to 

T,  the  osmotic  pressure  at  32°  should  be  51  x  -|—  =  54-1  cm.,  whilst 

2o7«2 

the  value  actually  found  was  54.6  cm. 

Finally,  it  remains  to  find  the  relationship  between  the  magnitudes 
of  the  gas  pressure  and  osmotic  pressure  under  comparable  condi- 
tions. This  can  also  be  done  on  the  basis  of  Pfeffer's  experiments 
with  cane  sugar,  for  example,  from  the  observation  that  at  o°  a  I  per 
cent,  solution  of  cane  sugar  at  o°  exerts  an  osmotic  pressure  equal  to 
49.3  cm.  of  mercury.  We  have  already  seen  that  the  molecular 
weight  of  a  gas  is  that  quantity  of  it  which,  when  present  in  22.4  litres 
at  o°,  exerts  a  pressure  of  I  atmosphere  (76  cm.).  Now  we  know 


196     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

from  chemical  considerations  that  the  molecular  weight  of  cane  sugar 
is  342.  As,  according  to  Pfeffer,  i  gram  in  100  c.c.  exerts  an  osmotic 
pressure  of  49.3  cm.  at  o°,  the  molecular  weight  of  cane  sugar,  342 
grams,  when  present  in  22.4  litres  or  22,400  c.c.,  must  exert  an  osmotic 
pressure  of 

49.3  x2fx-I22_o  =  75  Cm.  appro*. 

which  is,  within  the  limits  of  experimental  error,  the  same  as  the  gas 
pressure  it  would  exert  if  present  as  single  molecules  in  the  same 
volume  in  the  gaseous  form  in  the  absence  of  the  solvent.  That 
this  result  is  not  accidental  can  be  shown  by  similar  experiments  with 
other  substances  of  known  molecular  weight.  The  molecular  weight 
of  glucose  (grape  sugar)  is  180,  of  ethyl  alcohol  46,  and  it  has  been 
found  that  in  order  to  obtain  an  osmotic  pressure  of  I  atmosphere  at 
o°,  1 80  grams  of  grape  sugar,  or  46  grams  of  ethyl  alcohol,  must  be 
present  in  22,400  c.c.  of  solution.  From  this  and  other  considera- 
tions van  't  Hoffdrew  the  very  important  conclusion  that  the  osmotic 
pressure  exerted  by  any  substance  in  solution  is  the  same  as  it  would 
exert  if  present  as  gas  in  the  same  -volume  as  that  occupied  by  the  solu- 
tion, provided  the  solution  is  sufficiently  dilute. 

On  this  basis  is  founded  a  metWod  of  determining  molecular  weights 
in  solution  which  exactly  corresponds  with  that  used  for  gases  and 
vapours  (p.  1 10).  The  molecular  weight  of  a  dissolved  substance 
is  that  quantity  of  it  which,  when  present  in  22.4  litres  at  o°,  exerts  an 
osmotic  pressure  of  I  atmosphere. 

The  general  result  of  molecular  weight  determinations  in  solution 
by  this  method  is  that  the  values  obtained  usually  correspond  with 
the  simplest  formula  ascribed  to  a  substance  from  its  chemical 
behaviour.  Certain  important  exceptions  to  this  rule  will  be  con- 
sidered later. 

Determination  of  Molecular  Weights  in  Solution 
from  Depression  of  the  Freezing-point  and  Elevation 
of  the  Boiling-point  of  Solutions :— The  results  described  in 
the  last  section  may  be  summed  up  in  the  statement  that  the  osmotic 
pressure  of  solutions  containing  the  same  number  of  molecules  of 
different  solutes  in  equal  volumes  of  the  same  solvent  is  the  same  ;  in 
other  words,  the  osmotic  pressure  depends  on  the  number  and  not  on 
the  nature  of  the  particles.  The  determination  of  molecular  weights 
by  this  method  is,  however,  too  complicated  for  general  use,  and 
1  For  details  see  Physical  Chemistry,  pp.  109-138. 


OSMOTIC   PRESSURE  197 

it  has  been  found  much  more  convenient  to  measure  certain  other 
magnitudes  which  are  proportional  to  the  osmotic  pressure.  The  two 
most  important  of  these  are  the  depression  of  the  freezing-point  and 
the  elevation  of  the  boiling-point  produced  by  the  addition  of  a  known 
weight  of  a  soluble  substance  to  a  known  weight  or  volume  of 
solvent. 

It  has  been  known  for  a  long  time  that  a  dilute  solution  freezes  at 
a  lower  temperature  than  the  pure  solvent.  For  example,  a  solution 
containing  5  grams  of  sodium  chloride  in  100  grams  of  water  freezes 
at  -2.9°,  pure  ice  separating.  For  a  fixed  quantity  of  solvent  the 
lowering  of  the  freezing-point  is  proportional  to  the  amount  of  soluble 
substance  added,  and  for  a  definite  quantity  of  dissolved  substance 
the  depression  is  inversely  proportional  to  the  amount  of  solvent ;  in 
other  words,  the  depression  is  proportional  to  the  concentration  of  the 
solution.  Further,  like  the  osmotic  pressure,  the  depression  caused 
by  the  same  number  of  molecules  of  different  solutes  in  equal  amounts 
of  the  same  solvent  is  the  same.  The  lowering  produced  by  i  mol 
of  solute  in  100  grams  of  solvent  is  called  the  molecular  freezing-point 
depression ;  it  varies  with  the  nature  of  the  solvent.  The  molecular 
depression  K  for  a  solvent  is  found  once  for  all  by  experiment  with  a 
number  of  substances  of  known  molecular  weight,  and  the  molecular 
weight  of  any  other  substance  is  that  quantity  of  it  which,  when  dis- 
solved in  100  grams  of  the  solvent,  produces  the  depression  K.  The 
value  of  K  for  water  is  18.5°,  for  benzene  50°,  and  for  acetic  acid  39°. 

In  determining  molecular  weights  by  the  freezing-point  method,  it 
is  not  of  course  necessary  to  use  solutions  of  the  concentrations  stated 
above.  The  depression  is  determined  in  relatively  dilute  solutions, 
and  the  weight  in  grams  which  would  produce  a  depression  K  in  100 
grams  ot  solvent  is  calculated  on  the  assumption  that  even  in  concen- 
trated solutions  the  depression  is  proportional  to  the  concentration  of 
the  solution.  The  formula  for  calculating  the  results  is  readily  ob- 
tained as  follows :  If  g  grams  of  solute,  of  unknown  molecular 
weight  M,  dissolved  in  L  grams  of  solvent  (that  is,  100  gj~L  grams  in 
100  grams  of  solvent)  lowers  the  freezing-point  8  degrees,  whereas, 
according  to  the  above  law,  M  grams  of  solute  in  100  grams  of  solvent 
lowers  the  freezing-point  by  K  degrees,  we  have — 

:  8  : :  M  :  K  ;  whence  M  = 

In  order  to  illustrate  this  formula,  we  may  calculate  the  molecular 
weight  of  hydrogen  peroxide  from  the  observation  that  the  lowering 


198     A   TEXT-BOOK   OF    INORGANIC   CHEMISTRY 

of  the  free; 
pound  in  I 
formula — 


iyo       /l      lHi^l-J3VJVJJ\.     ^r      JLiNUJx \jrAINIL,     V^rlillVllO  I  Jx  I 

of  the  freezing-point  of  a  solution  containing  0.0943  grams  of  the  com- 
pound in  1 8.90  grams  of  water  is  0.270°.     Substituting  in  the  above 

r i_ 


ioo  x  0.0943  x  18.5 
M-       18.90x0.270      =34-3 

in  good  agreement  with  the  value,  34.0,  calculated  for  the  formula 
H202. 

The  effect  of  dissolved  substances  in  elevating  the  boiling-point  of 
a  solvent  exactly  corresponds  with  the  effect  in  depressing  the  freezing- 
point,  and  need  not  be  considered  in  detail.  In  the  former  case  also 
the  elevation  is  proportional  to  the  concentration  of  the  solution,  and 
the  same  number  of  molecules  of  different  substances  in  equal 
amounts  of  the  same  solvent  raise  the  boiling-point  of  the  solvent  to 
the  same  extent.  The  molecular  elevation  constant  can  be  found  by 
experiments  with  substances  of  known  molecular  weight  in  the  usual 
way.  The  value  for  water  is  5.2,  for  ethyl  alcohol  11.5,  and  for 
benzene  26.7. 

It  is  important  to  note  that  the  methods  of  determining  molecular 
weights  just  described  only  apply  when  the  pure  solvent,  unmixed 
with  the  solvent,  separates  (as  solid  or  vapour). 

The  magnitude  of  the  two  effects  just  discussed  is  much  less  than 
that  of  the  osmotic  pressure.  Thus  a  mol  of  solute  in  a  litre  of  water 
lowers  the  freezing-point  by  1.85°,  raises  the  boiling-point  by  0.52°, 
whereas  the  osmotic  pressure  of  the  solution  is  about  22.4  atmos- 
pheres. 

Eutectic  Mixtures.  Solid  Solutions — As  we  have  seen, 
the  lowering  of  the  freezing-point  of  a  fixed  quantity  of  solvent  by  a 
dissolved  substance  is  proportional  to  the  amount  of  the  latter  added. 
An  alternative  method  of  stating  the  facts  is  that  the  temperature  at 
which  solution  and  solid  solvent  are  in  equilibrium  is  the  lower  the 
greater  the  concentration  of  the  solution.  The  extent  to  which  the 
freezing-point  of  the  solvent  can  be  lowered  is  limited,  however,  by 
the  fact  that  the  solubility  of  solids  in  liquids  is  limited.  The  effect 
of  potassium  iodide  in  lowering  the  freezing-point  of  water  is  shown 
in  Fig.  47,  the  ordinates  representing  temperatures  and  the  abscissae 
concentrations.  The  point  A  represents  the  freezing-point  of  water 
(o°),  and  the  curve  AO  the  temperatures  at  which  ice  and  solution  are 
in  equilibrium  with  gradually  increasing  concentrations  of  salt.  O 
represents  the  point  at  which  the  solution  i&  saturated  with  potassium 
iodide  ;  at  this  time,  therefore,  solid  potassium  iodide,  as  well  as  ice^ 


OSMOTIC   PRESSURE 


199 


must  be  in  equilibrium  with  the  solution.  The  curve  BO  represents 
the  effect  of  temperature  on  the  solubility  of  potassium  iodide  ;  it  is 
the  curve  along  which  solid  potassium  iodide  is  in  equilibrium  with 
the  solution.  Since  the  solubility  diminishes  as  the  temperature  is 
lowered,  the  direction  of  the  curve  must  be  as  shown.  If  the  tem- 
perature of  a  saturated  solution  is  progressively  lowered,  salt  will 
separate,  and  the  concentration  will  fall  till  the  temperature  is  reached 
at  which  ice  also  begins  to  separate.  At  the  latter  temperature  salt 
and  ice  are  in  equilibrium  with  the  solution,  which  is  the  case  at  the 
point  O,  and  at  no  other  point.  The  latter  is  therefore  the  point  of 
intersection  of  the  freezing-point  and  solubility  curves.  If  a  mixture 
corresponding  with  the  composition  at  the  point  O  is  cooled  till  the 
freezing-point  is  reached, 
ice  and  salt  separate  in  the 
proportions  in  which  they 
are  present  in  the  mixture. 
Hence  the  composition  of 
the  solution  does  not  alter 
during  solidification,  and 
therefore  the  solution 
freezes  at  constant  tem- 
perature (p.  68)  like  a  pure 
substance.  The  mixture  of 
the  solid  components  in 
equilibrium  with  the  solu- 
tion is  termed  a  eutectic 
mixture;  when  one  of  the 
components  is  ice  it  is 
termed  a  cryohydrate.  The  latter  term  was  employed  because  it  was 
formerly  supposed  that  "cryohydrates"  were  definite  compounds  of 
salt  and  water,  but  it  can  be  shown  by  microscopic  examination  and 
in  other  ways  that  they  are  in  fact  mechanical  mixtures  of  the  two 
components. 

It  sometimes  happens  that  when  a  solution  is  cooled,  homogeneous 
crystals  containing  both  components  separate.  Mixtures  of  this  kind 
are  known  as  mixed' crystals  or  solid  solutions.  They  resemble  liquid 
solutions  in  being  homogeneous,  and  also  inasmuch  as  the  composi- 
tion varies  continuously  between  certain  limits.  The  freezing-point 
of  a  mixture  whose  components  form  solid  solutions  may  be  higher 
or  lower  than  those  of  the  components,  depending  upon  the  propor- 
tion in  which  the  latter  separate  from  solution. 


-3o 


Grams  KI    in  100  grams  solution 


30     40    50     60 
FIG.  47. 


70     80     90 


CHAPTER  XVI 

NITROGEN,  THE  ATMOSPHERE  AND  THE  ELEMENTS 
OF  THE  HELIUM  GROUP 

IT  has  already  been  stated  (p.  29)  that  the  atmosphere  is  composed 
essentially  of  the  two  gaseous  elements  oxygen  and  nitrogen. 
Within  the  last  twenty  years  it  has  been  discovered  that  no  less  than 
five  hitherto  unknown  elements  are  always  present  in  the  atmosphere, 
though  in  comparatively  small  amount.  Besides  these  elements,  the 
atmosphere  also  contains  varying  proportions  of  water  vapour,  carbon 
dioxide,  ammonia,  oxides  of  nitrogen,  and  various  other  substances. 
In  the  present  chapter  the  nitrogen  will  first  be  considered,  then  the 
atenosphere,  and  finally  the  preparation  and  properties  of  the  rare 
elements  just  referred  to. 

NITROGEN 

Symbol,  N.     Atomic  weight,  14.01.     Molecular  weight,  28.02. 

History — Nitrogen  was  first  isolated  by  Rutherford,  Professor 
of  Botany  at  Edinburgh,  in  1772.  He  kept  animals  in  a  confined 
volume  of  air  for  some  time,  and  after  removing  the  "fixed  air" 
(carbon  dioxide)  with  caustic  potash,  found  that  the  residual  air  was 
incapable  of  supporting  life  or  combustion.  He  termed  the  gas 
mephitic  air.  Lavoisier  was  the  first  to  regard  mephitic  air  as  an 
element.  He  termed  it  azote  (from  a,  privative,  and  £0077,  life)  in 
allusion  to  its  inability  to  support  life.  Chaptal  first  suggested  the 
name  nitrogen,  because  it  is  contained  in  nitre  or  saltpetre. 

The  opinion  held  up  to  1894,  that  atmospheric  nitrogen  is  a  single 
substance,  was  based  on  the  work  of  Cavendish  (cf.  p.  207). 

Occurrence — Nitrogen  occurs  free  in  the  atmosphere,  of  which 
it  constitutes  about  75.6  per  cent,  by  weight  or  78  per  cent,  by  volume. 
Free  nitrogen  also  occurs  to  a  very  small  extent  occluded  in  certain 
minerals.  In  combination  with  carbon  and  hydrogen  it  forms  an 
essential  constituent  of  plants  and  animals:-  In  combination  with 
hydrogen  it  forms  ammonia,  and  in  combination  with  oxygen  and 


NITROGEN  201 

other  elements  it  occurs  as  nitrates.  At  present  one  of  the  chief 
sources  of  combined  nitrogen  is  natural  sodium  nitrate  (Chili  saltpetre) 
Preparation — Nitrogen  can  be  obtained  by  two  principal 
methods :  (A)  by  removing  the  oxygen  from  purified  air  by  means  of 
easily  oxidizable  substances  ;  (B)  from  compounds  containing  nitrogen 
by  chemical  methods.  The  nitrogen  obtained  from  air  is  only  about 
99  per  cent,  pure,  as  it  is  mixed  with  argon  and  other  gases  of  the 
helium  group  which  cannot  easily  be  removed. 

(A)  The  methods  of  obtaining  nitrogen  by  removal  of  oxygen  from 
atmospheric  air  must  be  so  chosen  that  the  products  of  oxidation  can 
readily  be  separated  from  the  nitrogen.     The  more  important  are  as 
follows : — 

(1)  By  means  of  phosphorus.     A   small   piece   of  phosphorus  is 
caused  to  burn  in  air  confined  over  water  under  a  bell-jar.     Dense 
white  fumes  of  an  oxide  of  phosphorus  (phosphorus  pentoxide)  are 
produced,  which  dissolve  in  the  water  to  form  an  acid.     The  nitrogen 
obtained  in  this  way  is  impure,  because  the  phosphorus  becomes 
extinguished  before  the  oxygen  is  completely  removed. 

(2)  By  red-hot  copper.     The  air  is  first  freed  from  moisture  and 
carbon   dioxide    by    passing    through    U  -tubes   containing   suitable 
absorbing  agents,  and  is  then  drawn  through  a  long  tube  containing 
copper  filings  and  heated  in  a  furnace.     The  nitrogen  thus  prepared 
contains  argon,  but  is  otherwise  fairly  pure. 

(3)  In  the  last  method,  the  copper  slowly  becomes  inactive  through 
conversion  to  cupric  oxide.     If,  however,  a  mixture  of  air  and  ammonia 
gas  is  drawn  through  the  furnace  the  copper  oxide  is  continually 
reduced  to  metallic  copper  by  the  ammonia,  and  the  process  can  be 
continued  indefinitely.1 

(4)  By  alkaline  pyrogallol.    A  very  rapid  and  convenient  absorption 
of  oxygen  is  secured  by  shaking  air  with  a  solution  containing  pyro- 
gallol and  sodium  or  potassium  hydroxide.    The  oxygen  is  used  up  in 
oxidizing  th'e  pyrogallol  to  dark-coloured  complex  compounds.     This 
method  is  employed  for  estimating  oxygen  in  gaseous  mixtures. 

(B)  For  the  preparation  of  nitrogen  by  chemical  methods,  ammonium 
compounds  or  salts  of  nitrous  and  nitric  acids  are  generally  used. 

(5)  When  a  concentrated  aqueous  solution  of  ammonium  nitrite  is 
heated,  it  splits  up  directly  into  nitrogen  and  water: 

NH4NO,->N2  +  2H20. 

1  Equation:  3CuO  +  2NH3->3Cu  +  N2+3H2O.     The  excess  of  ammonia  can  be 
removed  by  passing  the  gas  through  water. 


202      A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

As   ammonium  nitrite   does   not  keep  well,  it  is  more   usual  to 
employ  a  mixture  of  sodium  nitrite  and  ammonium  chloride  : 


(6)  Nitrogen  is  also  obtained  by  heating  ammonium  bichromate 
(in   practice   a  mixture  of  ammonium   chloride  and  potassium  bi- 
chromate is  used)  : 

(N  H4)2Cr2O7->N2  +  Cr2O3  +  4H2O. 

(7)  By  passing  an  oxide  of  nitrogen,  for  example,  nitric  oxide,  NO, 
over  red-hot  copper  : 

2ND  +  2Cu->2CuO  +  N8. 

(8)  By  the  action  of  chlorine  on  ammonia,  the  latter  being  kept  in 
excess  : 


Physical  Properties  —  Nitrogen  is  a  colourless,  odourless,  taste- 
less gas.  The  density  of  pure  nitrogen  referred  to  air  is  0.9673,  whilst 
that  of  atmospheric  nitrogen  is  0.9721  (Rayleigh).  It  was  this  dis- 
crepancy in  the  densities  of  nitrogen  prepared  by  chemical  methods 
and  that  obtained  from  the  atmosphere  that  led  in  1894  to  the  discovery 
in  "atmospheric"  nitrogen  of  a  hitherto  unknown  element,  argon,  the 
density  of  which  is  about  .1.4  times  that  of  nitrogen.  Nitrogen  has 
been  obtained  both*  in  the  liquid  and  solid  form.  Liquid  nitrogen 
boils  at  -  196.6°  under  atmospheric  pressure;  its  critical  temperature 
is  -  146°  and  its  critical  pressure  35  atmospheres  (Olszewski).  Solid 
nitrogen  is  a  white  crystalline  substance  melting  at  -210°  (Fischer 
and  Alt)  ;  at  -252.5°  the  density  of  the  solid  is  1.0265  (Dewar). 

Nitrogen  is  only  very  slightly  soluble  in  water.  At  o°  i  c.c.  of 
water  dissolves  0.0239  c.c.,  at  20°  0.0164  c<c-}  anc^  a*  4°°  OtOIlS  c.c.  of 
nitrogen  at  76  cm.  pressure. 

Chemical  Properties  —  At  ordinary  temperatures  nitrogen  is 
a  very  inactive  element.  At  higher  temperatures  it  combines  directly 
with  a  number  of  other  elements,  more  particularly  lithium,  magnesium, 
barium,  calcium,  and  boron,  to  form  nitrides.  In  the  case  of  lithium, 
combination  takes  place  slowly  at  the  ordinary  temperature.  The 
formula  of  lithium  nitride  is  Li3N,  of  magnesium  nitride,  Mg3N2. 
Under  the  influence  of  the  electric  spark  or  the  electric  arc,  nitrogen 
and  hydrogen  form  small  quantities  of  ammonia,  NH3,  and  nitrogen 
and  oxygen  combine  to  form  a  brown  oxide  of  nitrogen,  NO2  (p.  227). 


THE   ATMOSPHERE  203 


THE  ATMOSPHERE 

As  already  mentioned,  the  atmosphere  is  mainly  composed  of 
oxygen  and  nitrogen,  in  the  proportion  of  I  part  of  oxygen  to  4 
parts  of  nitrogen  by  volume.  The  question  of  the  nature  of  the 
atmosphere  has  naturally  engaged  man's  attention  from  the  very 
earliest  times,  but  the  first  substantial  advances  in  knowledge  were 
made  in  connexion  with  the  phenomena  of  combustion.  Mayow, 
as  early  as  1674,  showed  that  only  a  portion  of  the  air  was  absorbed 
in  combustion  and  respiration,  and  similar  observations  were  made 
in  1692  by  Boyle.  The  complete  explanation  of  the  phenomena  of 
combustion,  and  the  clear  recognition  of  the  fact  that  air  is  essen- 
tially composed  of  an  elementary  substance  (oxygen)  which  combines 
with  other  substances  during  combustion,  and  of  another  element 
(nitrogen)  which  plays  no  part  in  combustion,  is  due  to  Lavoisier 
(p.  27). 

Composition  of  the  Atmosphere — The  relative  propor- 
tions of  oxygen  and  nitrogen  in  the  atmosphere  by  volume  can 
be  determined  by  absorbing  the  oxygen  from  a  known  volume  of 
air  by  one  of  the  methods  already  described  (p.  201),  and  measuring 
the  residual  nitrogen.  A  more  convenient  method  is  to  add  to  a 
measured  volume  of  air  in  a  eudiometer  excess  of  hydrogen,  explode 
the  mixture  by  means  of  an  electric  spark,  and  measure  the  residual 
gas.  As  one  volume  of  oxygen  combines  with  two  volumes  of 
hydrogen,  and  the  volume  of  the  water  formed  is  negligible,  $  of 
the  contraction  represents  the  volume  of  oxygen  in  the  original 
volume  of  air. 

The  relative  proportions  of  oxygen  and  nitrogen  by  weight  can  be 
determined  by  drawing  air,  freed  from  carbon  dioxide  and  moisture, 
over  heated  copper,  which  combines  with  the  oxygen  ;  the  nitrogen 
passes  on  into  a  previously  evacuated  flask  and  is  weighed.  The 
oxygen  is  determined  by  finding  the  increase  in  weight  of  the 
copper.  This  method  was  used  by  Dumas  and  Boussingault  (1841). 

The  result  of  numerous  analyses  of  air  from  the  most  varied 
sources  shows  that  there  are  slight  but  distinct  variations  in  its 
composition.  According  to  Angus  Smith,  the  proportion  of  oxygen 
in  towns,  especially  during  foggy  weather,  is  as  low  as  20.82  per  cent, 
by  volume,  whereas  on  open  ny>ors  and  mountains  it  is  as  high  as 
21.00.  As  a  mean  of  all  experiments,  it  has  been  found  tha^the 
atmosphere  contains  20.93  volumes  of  oxygen  and  79.07  volumes 


204     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

of  nitrogen  (including  argon),  or  23  per  cent,  of  oxygen  to  77  per 
cent,  of  nitrogen  by  weight.  A  litre  of  air  (free  from  carbon  dioxide 
and  moisture),  at  o°  and  76  cm.  pressure,  weighs  1.2933  grams. 

The  Atmosphere  a  Mixture — Owing  to  the  remarkable 
constancy  of  the  ratio  of  oxygen  and  nitrogen  in  the  atmosphere, 
it  has  sometimes  been  suggested  that  the  gases  are  in  chemical 
combination.  This  suggestion,  however,  is  untenable  for  many 
reasons : 

(1)  There  are  distinct  though  small  variations  in  the  composition 
of  the   air,  whereas   the  composition   of  a   chemical   compound   is 
constant. 

(2)  The  relative  amounts  of  the  two  gases  do  not  bear  any  simple 
relation  to  their  combining  weights. 

(3)  Heat  is  neither  given  out  nor  absorbed  when  the  gases  are 
mixed  in  the  proportions  in  which  they  are  present  in  air,  and  the 
resulting  mixture  has  all  the  properties  of  air. 

(4)  The  oxygen  and  nitrogen  retain  their  separate  properties  in 
air,  as   shown  (a)   by  the   fact  that   a   partial    separation   can    be 
effected  by  taking  advantage  of  their  different  rates  of  diffusion,  (b) 
when  air  is  shaken  up  with  water,  the  gases  dissolve  in  accordance 
with  their  respective  solubilities  and  partial  pressures  (p.  77), 

The  last  point  is  rather  interesting  as  it  indicates  a  comparatively 
simple  method  of  obtaining  a  mixture  rich  in  oxygen.  It  can  be 
calculated  that  the  air  expelled  from  water  by  boiling  contains  35 
volumes  of  oxygen  and  65  volumes  of  nitrogen. 

Substances  present  in  smaller  Proportion  in  the 
Atmosphere — Besides  oxygen,  nitrogen,  and  the  inactive  gases, 
the  atmosphere  contains  varying  proportions  of  aqueous  vapour, 
carbon  dioxide,  ammonia,  nitric  acid,  hydrogen,  and  a  number  of 
suspended  impurities,  including-  dust  particles,  bacteria  and  other 
organisms.  In  the  neighbourhood  of  towns,  sulphur  dioxide  and 
other  substances,  mainly  resulting  from  the  combustion  of  coal,  are 
found.  The  presence  of  ozone  in  normal  air  is  doubtful  (p.  138). 
Each  of  these  constituents  will  now  be  briefly  considered. 

Water  Vapour  —  At  a  given  temperature  a  definite  volume 
of  air  can  take  up  only  a  certain  amount  of  water  vapour;  when 
this  maximum  is  reached  the  air  is  said  to  be  saturated.  Air 
saturated  with  moisture  contains  at  o°  4.871  grams,  at  10°  9.362 
grams,  at  20°  17.157  grams,  and  att  30°  30.095  grams  of  aqueous 
vapmir  per  cubic  metre.  On  the  average,  the  amount  of  moisture 
is  about  two-thirds  of  the  saturation  value,  but  considerable  varia- 


THE   ATMOSPHERE  205 

tions  occur.  When  the  temperature  falls  below  that  at  which  the 
amount  of  moisture  is  just  sufficient  for  saturation,  the  excess  is 
deposited  as  dew,  mist  or  rain. 

The  amount  of  aqueous  vapour  in  the  atmosphere  can  be  deter- 
mined by  drawing  a  known  volume  of  air  through  bulbs  containing 
pumice  stone  soaked  with  concentrated  sulphuric  acid,  and  determin- 
ing the  increase  in  weight.  Hygrometers  are  also  used  for  the 
same  purpose. 

Carbon  Dioxide — The  normal  proportion  of  carbon  dioxide  in 
country  air  is  about  3  volumes  in  10,000,  but  in  large  towns,  where 
much  coal  is  burned,  it  may  rise  to  6  to  7  parts  in  10,000.  Air  contain- 
ing more  than  7  volumes  of  this  gas  in  10,000  is  considered  harmful 
for  respiration,  but  this  is  due  less  to  the  gas  itself  than  to  the 
accompanying  impurities. 

The  proportion  of  carbon  dioxide  in  the  atmosphere  is  continually 
increasing  owing  to  combustion  of  coal,  etc.,  and  to  respiration  (p.  335), 
but  it  is  simultaneously  being  taken  up  in  enormous  quantities  by 
plants  in  the  process  of  assimilation.  A  more  or  less  accurate 
compensation  is  thus  attained,  and  it  is  not  certain  whether  at 
the  present  time  the  amount  of  carbon  dioxide  is  increasing  or 
decreasing. 

The  proportion  of  carbon  dioxide  in  the  air  is  estimated  by  drawing 
a  measured  volume  of  it  through  a  solution  of  barium  hydroxide  of 
known  strength  : 

Ba(OH)2  +  C02->BaC03+H20, 

and  finding  by  volumetric  analysis  the  amount  of  unchanged  hydroxide. 

Ammonia  —  This  substance  results  from  the  decay  of  nitro- 
genous organic  matter  (p.*2i4).  Angus  Smith  found  0.5  to  i.o  gram 
in  10,000  grams  of  air  from  different  sources,  but  still  greater  varia- 
tions occur.  It  does  not  exist  to  any  great  extent  free  in  the  air, 
but  forms  salts  with  the  carbon  dioxide,  nitrous  and  nitric  acids 
which  are  also  present.  These  are  washed  into  the  soil  during 
rain,  and  form  a  very  important  source  of  the  combined  nitrogen 
required  for  the  growth  of  vegetation. 

Nitrous  and  Nitric  Acids — Under  the  influence  of  electrical 
discharges  in  the  atmosphere,  the  nitrogen  and  oxygen  com- 
bine to  form  oxides  of  nitrogen,  the  latter  then  reacting  with  water 
vapour  to  form  nitrous  and  nitric  acids.  As  already  mentioned, 
these  compounds  are  washed  into  the  soil  by  rain  and  form  a 
valuable  source  of  nitrogen  for  plants. 


206     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

It  was  formerly  supposed  that  plants  could  not  utilize  atmospheric 
nitrogen,  but  it  has  recently  been  shown  that  leguminous  plants, 
such  as  beans,  peas,  and  clovers,  can  make  use  of  nitrogen  by 
the  agency  of  bacteria.1  These  bacteria  are  to  be  found  in  nodules 
on  the  roots  of  the  plants  in  question,  and  by  means  not  yet  under- 
stood they  convert  the  atmospheric  nitrogen  (which  reaches  them 
through  the  pores  of  the  soil)  into  compounds  which  are  readily 
assimilated  by  plants. 

As  already  mentioned  (p.  31),  hydrogen  is  a  normal  constituent 
of  the  atmosphere,  but  the  proportion  is  very  small,  probably  not 
more  than  i  part  in  a  million. 

Suspended  Impurities  in  the  Atmosphere— The  air 
contains  a  large  amount  of  suspended  impurities,  which  are  rendered 
visible  when  a  ray  of  light  enters  a  darkened  room.  These  consist 
partly  of  inorganic  substances,  including  particles  of  salts,  etc.,  and 
partly  of  organic  substances,  including  bacteria  and  other  micro- 
organisms. Air  contains  under  normal  conditions  only  4  to  5 
micro-organisms  per  litre  on  the  average,  whereas  a  pure,  unfiltered 
river  water  contains  from  6,000  to  20,000  in  I  c.c.,  and  un- 
disturbed soil  about  100,000  in  i  c.c.  The  majority  of  the  micro- 
organisms in  the  air  consist  of  the  spores  of  yeast  and  of  moulds 
which  produce  fermentation.  The  bacteria  found  in  the  air,  in- 
cluding the  bacilli  of  various  diseases,  seldom  float  free  in  the  air, 
but  are  almost  always  attached  to  dust  particles,  and  the  latter, 
with  their  adherent  bacteria,  tend  to  settle  out  of  the  air  owing 
to  their  relatively  great  specific  gravity.  Bacteria  thus  reach  the 
air  chiefly  from  the  soil,  and  their  distribution  in  the  atmosphere  is 
favoured  by  dryness  of  the  soil  and  air  currents.  Air  can  be  almost 
entirely  freed  from  suspended  impurities*  by  drawing  it  through  a 
tube  loosely  packed  with  cotton  wool. 

The  dust  particles  in  the  air  act  as  nuclei  for  the  condensation  of 
moisture,  giving  rise  to  fogs.  This  may  be  well  illustrated  as  follows. 
A  large  flask  containing  a  little  water  at  the  bottom  is  partially  ex- 
hausted by  means  of  a  pump  ;  the  temperature  is  thus  considerably 
lowered,  and  a  fog  forms  in  the  flask.  If,  however,  the  air,  before 
being  admitted  to  the  flask,  is  filtered  through  cotton  wool,  no  fog  is 
produced  on  rapid  exhaustion. 

1  Pure  cultures  of  these  nitrifying  bacteria  (azotobakter)  have  been  placed  on 
the  market  under  the  name  of  nitragin  for  increasing  the  assimilation  of  nitrogen 
from  the  atmosphere.  The  results  hitherto  obtained  with  these  cultures  are 
rather  disappointing. 


THE   ELEMENTS   OF   THE   HELIUM    GROUP     207 
THE  ELEMENTS  OF  THE  HELIUM  GROUP 

Helium,  at.  wt.  =3.99.     Neon,  at.  wt.  =20.2.     Argon,  at.  wt.  =  39.88 
Krypton,  at.  wt.  =  82.9.     Xenon,  at.  wt.  =  130.2. 

History — In  a  paper  published  in  1785,  Cavendish  describes 
experiments  designed  to  find  whether  "  atmospheric  nitrogen "  is 
exclusively  composed  of  one  substance.  He  confined  a  mixture  of 
atmospheric  nitrogen  and  oxygen  over  potassium  hydroxide  solution 
and  passed  electric  sparks  through  the  mixture  till  no  further  absorp- 
tion took  place.  Only  a  small  bubble  remained,  and  Cavendish 
drew  the  conclusion  that  "  if  there  is  any  part  of  the  phlogisticated 
air  (nitrogen)  of  the  atmosphere  which  differs  from  the  rest,  and 
cannot  be  reduced  to  '  nitrous  acid,'  we  may  conclude  that  it  is  not 
more  than  ^%$  part  of  the  whole." 

In  1894  Lord  Rayleigh,  in  the  course  of  careful  experiments  on  the 
densities  of  gases,  found  that  whereas  a  litre  of  nitrogen  obtained 
from  chemical  compounds  weighs  1.2521  grams  under  normal  con- 
ditions, a  litre  of  nitrogen  prepared  from  the  atmosphere  weighs 
1.2572  grams.  The  experiments  undertaken  by  Rayleigh  and  by 
Ramsay  to  account  for  this  remarkable  difference  led  to  the  discovery 
of  argon,  a  gas  about  1.4  times  heavier  than  nitrogen.  It  is  an 
interesting  fact,  in  connexion  with  the  statement  of  Cavendish  just 
mentioned,  that  the  atmosphere  contains  rather  less  than  I  per  cent, 
by  volume  of  argon. 

While  searching  for  other  sources  of  argon,  Ramsay  (1895)  heated 
a  uranium-containing  mineral  called  cleveite,  and  obtained  a  gas  the 
spectrum  of  which  gave  a  line  (the  D3  line)  not  previously  observed 
with  any  terrestrial  gas,  but  which  had  been  detected  in  the  solar 
spectrum.  Lockyer  ascribed  this  characteristic  line  in  the  solar 
spectrum  to  an  element  which  he  called  helium,  and  this  name  was 
retained  by  Ramsay  when  the  element  was  discovered  on  the  earth. 

Argon  and  helium  differ  from  all  other  elements  in  being  remark- 
ably inactive  ;  so  far  it  has  not  been  found  possible  to  bring  them 
into  chemical  combination.  For  reasons  which  will  be  mentioned 
later,  Ramsay  was  led  to  search  for  other  inactive  gases,  and  in  1898 
three  other  gases,  belonging  to  the  same  family  as  argon  and  helium, 
were  discovered  by  Ramsay  and  Travers.  The  lightest  of  the  three, 
neon,  was  obtained  by  fractionating  a  large  quantity  of  argon,  the 
other  two,  krypton  (density  41.5)  and  xenon  (density  65.1),  were 
found  in  the  residues  after  the  evaporation  of  a  large  quantity  of 
liquid  air.  Each  of  these  gases  will  now  be  briefly  described. 


208     A 'TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

ARGON 

Symbol,  A.     Density,  19.94.     Atomic  and  molecular  weight,  39.9. 

Occurrence — Argon  is  present  in  the  atmosphere  to  the  extent  of 
0.94  per  cent,  by  volume,  or  1.3  per  cent,  by  weight.  It  is  also  found 
occluded  in  certain  minerals,  and  is  present  in  very  small  proportion 
in  most  natural  waters,  especially  in  those  from  certain  springs. 

Methods  of  Preparation — (i)  Argon  may  be  obtained  from 
the  atmosphere  by  passing  an  electric  discharge  through  a  mixture 
of  air  and  oxygen  in  the  presence  of  alkali  till  no  further  change  of 
volume  occurs  (p.  223).  The  excess  of  oxygen  is  then  removed  by 
alkaline  pyrogallate  and  the  residue  is  argon.  This  method  was  used 
by  Lord  Rayleigh  in  the  original  preparation  of  argon,  but  is  much 
less  convenient  than  that  now  to  be  mentioned. 

(2)  Argon  may  also  be  prepared  from  air  by  first  removing  the 
oxygen  with  red-hot  copper  and  then  passing  the  residual  gas,  freed 
from  carbon  dioxide  and  aqueous  vapour,  backwards  and  forwards 
over  heated  magnesium,  or  better,  over  heated  calcium  (a  mixture  of 
calcium  oxide  and  magnesium  powder  is  used)  till  no  further  diminu- 
tion of  volume  occurs.  The  magnesium  or  calcium  combine  with 
the  nitrogen  to  form  nitrides. 

The  argon  thus  obtained  is  contaminated  with  the  other  inactive 
gases,  but  can  be  purified  by  the  use  of  liquid  hydrogen  (Ramsay  and 
Travers).  When  the  crude  argon  is  passed  into  a  tube  immersed  in 
liquid  hydrogen,  helium  passes  on,  but  neon,  argon,  krypton  and 
xenon  are  liquefied  or  solidified.  On  allowing  the  mixture  to  warm 
up,  the  neon,  being  very  volatile,  first  distils  off,  then  argon  begins 
to  distil,  and  if  the  temperature  is  kept  at  a  suitable  point,  pure  argon 
distils  off,  while  the  less  volatile  krypton  and  xenon  remain  behind 
in  the  solid  form. 

Properties — Argon,  like  all  the  other  members  of  this  group,  is 
a  colourless,  odourless  gas.  Its  density  is  19.94,  and  therefore  its 
molecular  weight  is  39.88.  So  far,  all  attempts  to  make  argon  enter 
into  chemical  combination  have  been  fruitless,  so  that  its  atomic 
weight  cannot  be  determined  by  the  usual  volumetric  method  (p.  115). 
It  has,  however,  been  determined  by  the  following  method  that  argon 
and  the  other  gases  of  this  group  have  only  one  atom  in  the  molecule, 
so  that  the  atomic  and  molecular  weight  are  identical. 

It  can  be  shown1  that  the  molecular  heat2  of  a  monatomic  gas  at 

1  Physical  Chemistry,  p.  46. 

2  The  product  of  specific  heat  and  molecular  weight. 


THE   ELEMENTS    OF   THE    HELIUM    GROUP     209 

constant  volume,  that  is,  the  heat  required  to  raise  I  mol  of  the  gas 
i°  in  temperature,  is  3  calories,  but  more  heat  is  required  to  raise  a 
mol  of  a  polyatomic  gas  i°  in  temperature,  as  in  this  case  heat  is 
also  absorbed  in  doing  work  within  the  molecules.  Since  it  was  found 
by  experiment  that  the  molecular  heat  of  the  inactive  gases  was  3 
calories,  it  follows  that  they  are  monatomic.  The  molecular  heat  of 
mercury  vapour  is  also  3  calories  at  constant  volume,  and  the  con- 
clusion that  it  is  monatomic  has  been  confirmed  by  other  methods. 

Liquid  argon  boils  at  —  186°  and  solid  argon  melts  at  -  188°.  Its 
critical  temperature  is  -117.4°  and  the  critical  pressure  53  atmos- 
pheres. Argon  is  considerably  more  soluble  in  water*  than  nitrogen  ; 
at  15°  100  volumes  of  water  dissolve  about  4  volumes  of  argon. 

The  spectrum  of  argon  is  very  complex,  and  changes  markedly 
when  the  nature  of  the  discharge  is  altered.  With  the  intermittent 
discharge  the  glow  in  the  tube  is  red,  and  few  blue  lines  appear  in 
the  spectrum ;  with  an  oscillating  discharge  the  glow  is  bright  blue, 
the  red  lines  in  the  spectrum  disappear  or  become  faint,  and  many 
new  green  and  blue  lines  appear. 

HELIUM 

Symbol,  He.     Density,  2.0.     Atomic  and  molecular  weight,  3.99. 

Occurrence — Helium  occurs  in  certain  rare  minerals  containing 
uranium  or  thorium,  notably  in  cleveite,  uraninite,  monazite  sand,  and 
thorianite.  Onnes  obtained  the  200  litres  of  helium  used  in  the 
liquefaction  of  the  gas  by  heating  monazite  sand  (p.  493).  Ramsay 
found  that  thorianite,  a  rare  mineral  from  Ceylon  containing  76  to  78 
per  cent,  of  thorium  oxide,  along  with  the  oxides  of  uranium  and 
the  cerium  metals,  yielded  9  c.c.  of  helium  per  gram  when  heated  to 
redness.  Helium  is  also  present  in  certain  mineral  waters,  notably 
in  those  at  Bath,  which  contain  argon  with  8-10  per  cent,  of  helium. 
It  is  also  present  in  very  small  proportion,  according  to  Ramsay  to 
the  extent  of  0.0004  Per  cent,  by  volume,  in  the  atmosphere. 

Preparation — Helium  is  most  conveniently  obtained  from  the 
minerals  containing  it  ;  the  finely-powdered  material  is  heated  alone 
or  with  dilute  sulphuric  acid.  The  other  gases  generally  present  are 
removed  by  the  usual  methods. 

Helium   can   also   be   obtained    from   the   non-condensible   gases 

obtained  in  the  process  of  making  liquid  air.     The  gases  escaping 

from  the  liquefier  contain  a  considerable  proportion  of  helium  and 

neon.      All  the   gases   except   helium   are  condensed  and  in  fact 

14 


210    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

solidified  at  the  temperature  of  liquid  hydrogen,  so  that  the  helium 
can  be  pumped  off  practically  pure. 

Properties — Next  to  hydrogen,  helium  is  the  lightest  gas 
known,  its  density  being  only  2.  After  many  failures,  it  was  liquefied 
for  the  first  time  by  Kammerlingh  Onnes  of  Leyden  in  1908.  200 
litres  of  the  gas,  cooled  by  the  evaporation  of  liquid  hydrogen  under 
reduced  pressure,  were  circulated  round  the  liquefier  (of  the  Linde 
type)  for  about  three  hours,  and  finally  about  60  c.c.  of  liquid  helium 
was  obtained.  The  liquid  boiled  about  4.3°  abs.  (-268.7°  C.)  and 
its  density  was  0.15;  the  critical  temperature  is  about  5°  abs.  By 
evaporation  of  liquid  helium  a  temperature  within  3°  of  the  absolute 
zero  was  reached. 

Helium  is  only  very  slightly  soluble  in  water  ;  at  o°  100  c.c.  of 
water  dissolve  0.015  c-c-j  and  at  20°  0.0138  c.c.  of  the  gas. 

With  the  intermittent  discharge  at  7  to  8  mm.  pressure,  helium  gives 
a  yellow  spectrum,  the  characteristic  so-called  D3  line  (p.  207)  attain- 
ing its  greatest  intensity  ;  when  the  pressure  is  reduced  to  i  to  2  mm. 
the  tube  emits  a  brilliant  green  light.  The  yellow  D3  line,  a  red 
line,  and  two  lines  in  the  green  are  most  characteristic. 

NEON 
Symbol,  Ne.     Density,  10.1.     Atomic  and  molecular  weight,  20.2. 

Neon  was  isolated  from  the  less  condensible  portion  of  liquid  air, 
as  already  mentioned  under  helium.  The  complete  separation  from 
argon  and  helium  was  rendered  possible  by  the  preparation  of  liquid 
hydrogen  in  quantity  by  Travers.  The  argon  and  neon  were  solidified 
in  a  bulb  immersed  in  liquid  hydrogen  and  the  helium  pumped  off ; 
the  temperature  was  then  allowed  slowly  to  rise.  As  neon  boils  at  a 
lower  temperature  than  argon,  the  first  fraction  was  rich  in  neon 
and  contained  very  little  argon. 

Properties — The  boiling-point  and  melting-point  of  neon  have 
not  been  accurately  determined.  When  an  electric  discharge  is 
passed  through  it  in  a  vacuum  tube  the  colour  is  orange  red.  The 
spectrum  contains  a  number  of  lines  in  the  red,  a  prominent  line,  D6, 
in  the  yellow,  and  a  number  of  strong  green  lines. 

KRYPTON  AND  XENON 

Krypton,  Symbol  Kr.     Density,  41.5.      Atomic  weight,  83.0. 
Xenon,  Symbol  X.          Density,  65.1.       Atomic  weight,  130.2. 

Preparation — These  gases  were  isolated  from  the  residue 
obtained  in  the  evaporation  of  a  large  quantity  of  liquid  air.  The 


THE   ELEMENTS   OF  THE   HELIUM   GROUP     211 

oxygen  and  nitrogen  were  removed  by  the  usual  methods,  the  residue, 
containing  krypton  and  xenon,  was  again  liquefied  and  the  substances 
separated  by  fractional  distillation,  krypton  being  considerably  more 
volatile  than  xenon. 

Properties — Krypton  boils  about  -152°  C,  and  the  solid 
melts  at  —  169°  ;  xenon  boils  about  —  107°  C.  and  the  solid  melts 
about  -  140°. 

Like  helium  and  neon,  the  spectrum  of  krypton  is  independent  of 
the  nature  of  the  electric  discharge.  The  spectrum  contains  a  few 
distinct  lines  in  the  red,  but  a  yellow  and  a  green  line  are  the  most 
brilliant.  The  green  line  appears  to  be  identical  with  a  prominent 
line  in  the  spectrum  of  the  aurora  borealis. 

The  spectrum  of  xenon,  like  that  of  argon,  varies  with  the  nature 
of  the  discharge.  With  the  intermittent  discharge  the  glow  is  blue  ; 
when  a  Leyden  jar  and  spark  gap  are  introduced  in  the  circuit,  the 
colour  is  green. 

Summary  of  Group — The  members  of  this  group,  like  the 
halogens,  form  a  family  of  elements,  whose  properties  vary  regularly 
with  increasing  atomic  weight.  The  more  important  physical  pro- 
perties of  the  elements  are  summarized  in  the  table  : — 


Helium. 

Neon. 

Argon. 

Krypton. 

Xenon. 

Atomic  weight 
Boiling-point  . 

3-99 
-268.7° 

20.2 
-233° 

39-88 
-1  86° 

82.9 
-151.7 

130.2 
—  106.9° 

Melting-point  . 
Critical  temperature 

-268° 

-i  88° 
-117.4° 

-I69° 
-    62.5° 

—  140° 
+   16.6 

Atomic  volume 

... 

... 

32-9 

38.5 

42.7 

The  elements  of  this  group  resemble  each  other  in  being  chemically 
inactive ;  so  far  none  of  them  have  been  brought  into  chemical 
reaction.  We  may  therefore  say  that  the  valency  of  the  group  is 
zero.  They  are  all  monatomic,  as  shown  by  considerations  based  on 
their  specific  heat. 

It  may  be  stated  that  a  number  of  attempts  have  been  made  to 
obtain  other  elements  belonging  to  this  group,  but  hitherto  without 
definite  success.  Ramsay  and  Moore  examined  the  residues  from 
the  evaporation  of  over  100  tons  of  liquid  air  made  by  Claude's 
method  (p.  74),  but  found  no  element  denser  than  xenon. 


A  TP:XT-BOOK  OF  INORGANIC  CHEMISTRY 


The  proportion  of  these  gases  present  under  normal  circumstances 
in  the  atmosphere  is  given  in  the  following  table  (Ramsay)  : — 


Parts  by  Weight 
in  i  Part  of  Air. 

Parts  by  Volume 
in  i  Part  of  Air. 

0.00000056 

0.000004 

0.0000086 

0.0000123 

0.0136 

o.ooq^7 

0.00028 

O  00009 

Xenon     

0.00005 

O.OOOOI 

CHAPTER  XVII 

COMPOUNDS   OF    NITROGEN   WITH    HYDROGEN 
AND   WITH    THE    HALOGENS 

'"pHREE  compounds  of  nitrogen  and  hydrogen  are  known.     Theii 
A       names  and  formulae  are  as  follows  : — 

Ammonia       .         .        .        .        .         .     NH3 

Hydrazine N2H4  or  NH2'NH2 

Hydrazoic  acid  or  azoimide  .         .         .     N3H 

It  will  also  be  convenient  to  describe  in  this  chapter  a  compound 
of  nitrogen  with  hydrogen  and  oxygen. 

Hydroxylamine      .  ....     NH2'OH 

AMMONIA 

Formula,  NH3.     Molecular  weight,  17.03. 

History — Compounds  of  ammonia  have  been  known  from  the 
earliest  times.  Ammonium  chloride  was  formerly  imported  into 
Europe  from  Egypt,  where  it  was  prepared  (by  sublimation)  from  the 
soot  obtained  on  burning  camel's  dung,  which  was  used  as  fuel  near 
the  temple  of  Jupiter  Ammon  in  Libya.  From  this  circumstance 
the  salt  was  called  sal-ammoniac,  whence  the  name  ammonia  is 
derived.  At  a  later  period  compounds  of  ammonia  were  prepared  by 
the  dry  distillation  of  horns,  hoofs  and  other  animal  matter.  From 
this  mode  of  preparation  is  derived  the  name  spirits  of  hartshorn, 
sometimes  applied  to  ammonia.  Ammonia  gas  was  first  obtained  by 
Priestley  (1774),  who  collected  it  over  mercury. 

Occurrence — As  already  mentioned,  ammonia  occurs  in  the 
atmosphere,  partly  in  the  free  starte,  but  chiefly  as  carbonate  and 
nitrate.  It  is  found  as  the  chloride  and  sulphate  near  active  vol- 
canoes. As  it  is  a  product  of  the  decay  of  animal  and  vegetable 
matter,  it  is  found  in  the  combined  state  in  soils  and  in  natural 
waters. 


2i4    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Preparation  —  (i)  Traces  of  ammonia  are  obtained  by  passing 
a  silent  electrical  discharge  through  a  mixture  of  nitrogen  and 
hydrogen  : 


As  the  reaction  is  a  reversible  one,  and  the  equilibrium  lies  very 
near  the  left-hand  side,  only  a  small  proportion  of  ammonia  is  formed, 
but  if  it  is  removed  by  an  acid  as  fast  as  it  is  produced  the  reaction 
goes  completely  in  the  direction  of  the  upper  arrow.  This  is  an 
excellent  illustration  of  the  influence  of  the  concentration  of  the 
reacting  substances  on  the  equilibrium  (p.  165). 

(2)  By  the  dry  distillation   of   organic    nitrogenous    compounds 
(animal  and  vegetable  matter).     When  coal,  which  results  from  the 
slow  decay  of  vegetation,  is  heated  out  of  contact  with  air,  ammonia 
is  one  of  the  products.     It  is  absorbed  in  water,  and  the  solution,  the 
so-called  "  ammoniacal  liquor  of  the  gas-works,"  is   the  chief  com- 
mercial source  of  ammonia  compounds.     The  ammoniacal  liquor  is 
impure,  but  a  fairly  pure  ammonium  salt  can  be  obtained  by  heating 
the  liquor  with  slaked  lime,  and  absorbing  the  ammonia  in  hydro- 
chloric or  sulphuric  acid  (cf.  p.  418). 

(3)  When  organic  nitrogenous  compounds  are  heated  for  some 
hours  with  concentrated  sulphuric  acid,  the  nitrogen  is  completely 
converted  into  ammonia,  and  remains  in  the  solution  as  ammonium 
sulphate.      Kjeldahl's  method   for  estimating   nitrogen    in  organic 
compounds  is  based  on  this  reaction.     The  nitrogen  of  many  organic 
compounds  is  also  converted  into  ammonia  on  heating  with  a  con- 
centrated solution  of  an  alkali,  especially  in  the  presence  of  potassium 
permanganate. 

(4)  Nitrates  and  nitrites  are  reduced  to  ammonia  by  "nascent" 
hydrogen,  that  is,  by  hydrogen  generated  in  the  solution  : 


2O-|-NH3. 

This  method  is  made  use  of  in  the  estimation  of  nitrites  and  nitrates 
in  drinking  water. 

(5)  Pure  ammonia  gas  is  best  prepared  by  heating  solid  ammonium 
chloride  with  dry  calcium  hydroxide.  Ammonium  salts  contain  the 
univalent  group  NH4,  which  plays  the  part  of  a  univalent  metal 
(p.  418).  The  formula  for  ammonium  chloride  is  therefore  NH4C1, 
and  the  reaction  just  described  may  be  written  as  follows  :  — 

2NH4Cl+Ca(OH)2-»2NH8t 


NITROGEN-HYDROGEN   COMPOUNDS          215 

The  gas  is  dried  by  passing  it  through  a  long  tube  containing 
lumps  of  calcium  oxide,  and  can  be  collected  over  mercury  or  by 
downward  displacement  of  air. 

Physical  Properties  —  Ammonia  is  a  colourless  gas,  with  a 
pungent  smell  and  caustic  taste.  Its  density  referred  to  air  is  0.597. 
It  can  be  condensed  to  a  colourless  liquid,  which  boils  at  -32.5°, 
and  on  further  cooling  yields  a  solid  which  melts  at  —77°.  Its 
critical  temperature  is  131°  and  critical  pressure  113  atmospheres. 
Liquefied  ammonia  is  an  excellent  solvent.  It  absorbs  a  large 
amount  of  heat  on  evaporation,  a  fact  taken  advantage  of  in  machines 
for  making  artificial  ice. 

Ammonia  gas  is  very  soluble  in  water.  At  o°  i  c.c.  of  water 
absorbs  1299  c-c«?  at  I0°  7^3  c-c-»  anc*  at  28°  595  c.c.  of  ammonia, 
measured  at  o°  and  760  mm.  pressure  (Raoult).  The  gas  is  com- 
pletely expelled  from  its  aqueous  solution  by  boiling. 

Quite  recently  the  equilibrium  2NH3^tN24-3H2  has  been 
thoroughly  investigated.  The  equilibrium  mixture  at  700°  and 
30  atmospheres  pressure  contains  less  than  I  per  cent,  of  ammonia  ; 
at  ordinary  pressures  the  proportion  is  of  course  much  smaller.  As 
the  heat  of  formation  of  ammonia  from  its  elements  is  positive  (11,890 
cal.  for  17  grams)  the  equilibrium  is  displaced  in  favour  of  the  ammonia 
as  the  temperature  is  lowered  (p.  172). 

Chemical  Properties  —  The  most  striking  chemical  property 
of  ammonia  is  that  it  combines  directly  with  acids  to  form  salts, 
thus~  NH3  +  HC1=NH4C1 

2NH3+H2SO4  =  (NH4)2SO4. 

As  already  mentioned,  the  salts  obtained  in  this  way  are  termed 
ammonium  salts,  the  NH4,  or  ammonium  group,  playing  the  part  of 
a  metal.  Ammonium  salts  can  be  vaporized  by  heating,  and  the 
densities  of  the  vapours  are  considerably  less  than  the  values  cal- 
culated from  their  formulas.  In  the  case  of  ammonium  chloride,  for 
instance,  the  molecular  weight  calculated  from  its  formula  is  53.5,  but 
the  density,  instead  of  being  26.7,  is  only  about  half  that  value.  The 
simplest  explanation  is  that  in  the  form  of  vapour  ammonium  chloride 
is  dissociated  almost  completely  into  ammonia  and  hydrogen  chloride, 
according  to  the  reversible  equation 


This  suggestion  has  been  fully  confirmed  by  experiment,  a  partial 
separation  of  the  products  of  dissociation  having  been  effected  by 
taking  advantage  of  their  different  rates  of  diffusion. 


216     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 


The  conclusion  which  may  be  drawn  from  the  above  results,  that 
ammonia  is  basic  in  character,  is  confirmed  by  the  observation  that 
its  aqueous  solution  is  alkaline  to  litmus.  It  has  already  been  stated 
that  all  bases  in  solution  contain  the  OH  group,  and  we  may  there- 
fore assume  that  ammonia  reacts  with  water  to  form  ammonium 
hydroxide,  NH4OH,  corresponding  with  the  alkali  hydroxides  KOH 
and  NaOH. 

It  is  probable  that  when  ammonia  unites  with  hydrogen  chloride, 
water,  etc?,  the  nitrogen  changes  from  the 
trivalent  to  the  pentavalent  state  (p.  578). 

Ammonia  does  not  burn  in  air,  but  in  oxygen 
it  burns  with  a  yellow  flame.  The  chief  pro- 
ducts are  water  and  nitrogen,  but  traces  of 
oxides  of  nitrogen  are  also  formed.  This  in- 
teresting reaction  is  best  shown  by  means  of 
the  arrangement  represented  in  Fig.  48.  *  Am- 
monia obtained  by  boiling  the  aqueous  solution 
in  the  flask  passes  through  a  glass  tube 
enclosed  in  a  wider  tube.  When  oxygen  is 
admitted  to  the  cylindrical  space  between  the 
two  tubes,  the  ammonia  can  be  burned  at  the 
end  of  the  narrow  tube  ;  but  if  the  supply  of 
oxygen  is  cut  off,  the  flame  becomes  smaller 
and  finally  goes  out. 

When  a  mixture  of  ammonia  and  air  is 
passed  over  platinum  black,  heated  to  redness, 
nitric  acid  is  obtained  (Ostwald),  a  process 
which  has  recently  become  of  great  com- 
mercial importance. 

-  3      Ammonia  is  completely  decomposed  by  free 
FIG-  48'  chlorine  : 


When  excess  of  ammonia  is  present,  the  hydrochloric  acid  is  con- 
verted into  ammonium  chloride.  Hypochlorites,  e.g.  solution  of 
chlorinated  lime,  also  oxidize  ammonia  with  liberation  of  nitrogen. 

Many  metals  decompose  ammonia  into  its  elements  at  high  tem- 
peratures, and  in  this  way  a  number  of  metallic  nitrides  have  been 
obtained.  Nitrides  may  in  fact  be  regarded  as  being  derived  from 
ammonia,  by  displacing  the  hydrogen  by  metals.  They  are  decom- 
posed by  water,  forming  ammonia  and  metallic  hydroxides  : 


NITROGEN-HYDROGEN  COMPOUNDS 


217 


Just  as  water  unites  with  a  number  of  metallic  salts  as  so-called  water 
of  crystallization,  so  many  solid  compounds  of  metallic  salts  with 
ammonia  are  known.  Of  these,  the  compounds  with  silver  chloride, 
2AgCl,3NH3  and  AgCl,3NH3,  and  the  compounds  with  copper  sul- 
phate, CuSO4,4NH3;  CuSO4,2NH3  and  CuSO4,NH3, are  most  familiar. 

Ammonia  can  be  liberated  from  its  compounds  by  warming  with  an 
alkali,  and  is  detected  by  its  characteristic  smell,  and  by  the  fact  that 
it  turns  litmus  paper  blue.  Traces  of  ammonia,  or  of  ammonium 
compounds,  are  detected  by  Nessler's  reagent  (p.  459). 

Composition  of  Ammonia — If  the  formula  of  ammonia  is 
NH3,  the  equation  representing  its  formation  from  nitrogen  and 
hydrogen  must  be  as  follows  : — 

N2     +     3H2     =     2NH3 


i  vol. 


2  VOls. 


3  vols. 

that  is,  one  volume  of  nitrogen  and  three  volumes  of  hydrogen  unite 
to  form  two  volumes  of  ammonia.  Perhaps  the  simplest  method  of 
confirming  the  formula  is  to  submit  a  measured  volume  of  the  gas  to 
electric  sparks  in  the  apparatus  represented  on  p.  223  till  its  volume 
no  longer  changes.  On  again  adjusting  to  atmospheric  pressure,  it 
will  be  observed  that  the  volume  has  practically  doubled,  in  accord- 
ance with  the  above  equation.  Oxygen  is  then  added  in  excess,  and 
an  electric  spark  passed  through  the  mixture.  From  the  contraction, 
due  to  the  formation  of  water,  the  volume  of  hydrogen  in  the  original 
mixture  can  be  calculated,  and  it  may  be  shown  that  one 
volume  of  nitrogen  to  three  of  hydrogen  were  present. 

The  fact  that  ammonia  is  decomposed  by  chlorine  may 
also  be  taken  advantage  of  to  show  the  composition  of 
the  former  gas.  A  long  glass  tube,  provided  with  a  stop- 
cock and  a  funnel  above,  is  marked  off  externally  into 
three  equal  sections  by  indiarubber  rings  (Fig.  49),  and 
filled  with  chlorine  by  displacement  of  a  strong  solution 
of  brine.  A  concentrated  solution  of  ammonia  is  poured 
into  the  funnel,  and  then  admitted,  a  drop  at  a  time, 
into  the  chlorine  by  cautiously  turning  the  tap.  At  first 
a  yellowish  flame  follows  the  admission  of  each  drop, 
but  after  a  time  this  effect  ceases,  as  all  the  chlorine  is 
decomposed.  When  a  considerable  excess  (three  or  four 
c.c.)  of  ammonia  solution  has  been  added,  so  as  to  ensure 
complete  conversion  into  ammonium  chloride  and  nitrogen  (p.  220), 
excess  of  dilute  sulphuric  acid  is  added  through  the  tap  to  neutralize 


FIG.  49.  • 


218     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

the  ammonia,  and  water  is  then  cautiously  drawn  into  the  tube  in  the 
same  way,  until  no  more  will  enter.  If  the  experiment  has  been 
performed  in  such  a  way  that  no  gas  has  escaped  through  the  tap,  the 
residual  gas,  which  is  practically  under  atmospheric  pressure,  fills  one 
of  the  divisions  of  the  tube,  and  may  easily  be  shown  to  be  nitrogen. 
As  hydrogen  and  chlorine  combine  in  equal  volumes,  the  original 
three  volumes  of  chlorine  must  have  combined  with  three  volumes  of 
hydrogen,  and  one  volume  of  nitrogen  has  been  liberated.  It  follows 
that  ammonia  is  formed  by  the  combination  of  one  volume  of 
nitrogen  with  three  volumes  of  hydrogen.  Hence  the  formula  is 
(NH3)*,  and  as  the  molecular  weight  is  17,  it  follows  that  the  molecular 
formula  for  ammonia  is  NH3. 

The  same  fact  may  also  be  proved  by  electrolyzing  a  concentrated 
aqueous  solution  of  ammonia,  to  which  some  ammonium  sulphate  has 
been  added,  in  the  apparatus  represented  on  p.  14.  Hydrogen  is 
liberated  at  the  negative  and  nitrogen  at  the  positive  pole,  and  the 
volume  of  the  former  gas  is  three  times  that  of  the  latter. 

HYDRAZINE,  N2H4 

Preparation—  Hydrazine,  or  diamide,  H2N  —  NH2,  was  discovered  by  Curtius 
(1887).  It  may  be  obtained  by  the  reduction  of  hyponitrous  acid  (p.  234)  : 


but  is  most  easily  prepared  by  adding  sodium  hypochlorite  to  excess  of  a  con. 
centrated  solution  of  ammonium  hydroxide,  which  contains  about  0.02  per  cent. 
of  glue.  The  excess  of  ammonia  is  removed  by  boiling,  the  solution  is  concen- 
trated and  on  addition  of  sulphuric  acid  hydrazine  sulphate,  N2H4,H2SO4, 
crystallizes  out  (Raschig).  It  is  probable  that  monochloramide,  NH2C1,  is  an 
intermediate  product,  and  is  acted  on  by  ammonia  to  form  hydrazine: 

NH2C1  +  NH3->NH2-NH3-HC1. 

Hydrazine  Hydrate,  N2H4,H2O,  is  obtained  by  distilling  a  salt  of  hydrazine, 
for  instance,  the  sulphate  or  bromide,  with  a  very  concentrated  solution  of  potas- 
sium hydroxide  under  reduced  pressure.  It  is  a  colourless  liquid,  which  boils  at 
120°,  smells  of  ammonia,  and  has  a  powerful  corrosive  action.  It  can  be  mani- 
pulated only  in  vessels  of  silver  or  platinum.  At  100*  it  is  partially  dissociated 
into  hydrazine  and  water  ;  the  dissociation  is  practically  complete  at  [140°. 
Hydrazine  hydrate  was  known  for  some  years  before  hydrazine  itself  had  been 
obtained.  The  latter  is  prepared  from  the  hydrate  by  distilling  with  excess  of 
barium  oxide  under  reduced  pressure  : 

•  N2H4-H20  +  BaO  =  N3H4+Ba(OH)2. 

Properties  —  At  ordinary  temperatures  hydrazine  is  a  colourless  liquid  which 
boils  at  113°,  and  can  easily  be  obtained  in  the  form  of  a  solid  melting  at  i°. 
The  density  of  the  liquid  at  15°  is  1.0114.  Like  ammonia,  hydrazine  is  a  base, 


NITROGEN-HYDROGEN  COMPOUNDS  219 

combining  with  acids  to  form  salts.  One  molecule  is  able  to  combine  with  one 
and  with  two  molecules  of  hydrochloric  acid,  forming  the  compounds  N2H4,HC1 
and  N2H4,2HC1.  It  combines  very  readily  with  water  to  form  the  hydrate 
N2H4,H20. 

Hydrazine  and  its  salts  are  very  powerful  reducing  agents,  owing  to  the  readi- 
ness with  which  the  base  is  oxidized  to  water  and  nitrogen  : 

N2H4  +  H20->N2 + 2H2O. 

Thus  from  the  salts  of  the  noble  metals  (gold,  mercury,  etc. )  the  free  metals  are 
precipitated  in  the  cold,  and  cupric  salts  are  readily  reduced  to  cuprous  oxide  and 
ferric  salts  to  the  ferrous  state.  With  the  halogens  free  nitrogen  and  the 
corresponding  halogen  acids  are  obtained.  In  the  presence  of  finely  divided 
platinum  hydrazine  in  aqueous  solution  breaks  up  into  nitrogen,  hydrogen,  and 
ammonia,  the  ratio  in  which  the  products  are  formed  depending  on  the  degree  of 
acidity  or  alkalinity  of  the  solution  (Tanatur). 

Although  the  formula  of  the  hydrate  is  usually  written  N2H4,H2O,  and  that 
of  the  hydrochloride  as  N2H4,HC1,  it  is  probable,  from  analogy  with  the  corre- 
sponding ammonium  compounds,  that  one  of  the  nitrogens  in  the  former  com- 
pounds is  quinquevalent,  and  hence  that  the  formula  of  the  hydroxide  is  as  follows  : 

/H 
rl\  /  TT 

>N-N<g 
H/         \OH 

The  relatively  great  stability  of  hydrazine  hydrate  as  compared  with  ammonium 
hydroxide  is  worthy  of  note. 

HYDRAZOIC  ACID  OR  AZOIMIDE,  N3H 

Preparation— This  remarkable  compound  was  first  obtained  by  Curtius  in  1890. 
It  is  most  conveniently  prepared  by  passing  nitrous  oxide,  N2O,  over  fused 
sodium  amide,  NaNH2,  at  190°  and  distilling  the  resulting  sodium  hydrazoate 
(sodium  azide)  with  dilute  sulphuric  acid : 

N\          H\  N\ 

||    >O+       /N-Na-M|    / 
N/         H/  N/ 

2NaN3+  H2SO4->2N3H  +  Na2SO4. 

Hydrazoic  acid  can  also  be  obtained  by  shaking  an  aqueous  solution  of  hydrazine 
with  a  3  per  cent,  solution  of  nitrogen  trichloride  in  benzene,  the  liquid  being  kept 
alkaline  during  the  reaction  by  addition  of  sodium  hydroxide  (Tanatur).  A  third 
method  of  preparation  is  to  mix  in  the  cold  a  solution  of  hydrazine  sulphate 
(5  grams)  and  potassium  nitrite  (3.3  grams  in  200  c.c.  water) ;  after  the  evolution 
of  gas  has  ceased  the  product  is  distilled : 

NH2  N\ 

|        +NO-OH->|| 
NH2  N 

Properties — By  fractional  distillation  of  the  aqueous  solution  a  product  con- 
taining 91  per  cent,  of  hydrazoic  acid  is  obtained,  which  yields  the  anhydrous 


220     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

acid  on  treatment  with  calcium  chloride.  The  anhydrous  acid  is  a  colourless, 
mobile  liquid,  which  boils  at  37°,  has  a  penetrating  odour,  and  readily  decom- 
poses explosively  into  its  elements  :  2NSH—  >3N2  +  H2. 

Hydrazoic  acid  is  a  monobasic  acid  of  medium  strength  and  readily  dissolves 
metals  such  as  zinc  and  tin  with  formation  of  salts,  the  azides.  The  azides  of 
heavy  metals,  e.g.  lead  azide,  PbN6,  and  silver  azide,  AgN3,  are  highly  explosive. 

A  comparison  of  the  three  compounds  of  nitrogen  and  hydrogen  just  considered 
shows  that  the  basic  character  diminishes  as  the  proportion  of  nitrogen  increases. 

HYDROXYLAMINE,  NH2OH 

Preparation—  Hydroxylamine,  discovered  by  Lessen  in  1865,  can  be  obtained 
by  the  action  of  "nascent"  hydrogen  on  nitric  oxide  or  nitric  acid  : 


The  most  convenient  method  is  to  pass  nitric  oxide  through  a  mixture  in  which 
hydrogen  is  being  generated  by  the  action  of  hydrochloric  acid  on  tin,  hydroxyl- 
amine  hydrochloride,  NH2OH'HC1,  being  thus  obtained.  The  tin  is  removed 
by  means  of  hydrogen  sulphide,  and  the  hydroxylamine  hydrochloride  dissolved 
out  with  absolute  alcohol. 

Hydroxylamine  can  be  obtained  by  reducing  nitric  acid  by  means  of  hydrogen 
generated  electrolytically.  By  using  a  cathode  of  amalgamated  lead  (lead  coated 
with  mercury),  80  per  cent,  of  the  nitric  acid  used  can  be  converted  into  hydroxyl- 
amine. 

Anhydrous  hydroxylamine  can  be  obtained  most  readily  by  heating  hydroxyl- 
amine phosphate  in  a  vacuum,  the  free  base  passing  over  at  135°  to  137°  : 
(NH2OH)3-H3P04->3HN2OHt  +  H3PO4. 

Properties  —  Hydroxylamine  occurs  in  colourless  needles  which  melt  at  33°. 
The  liquid  boils  without  decomposition  at  56°  to  58°  under  22  mm.  pressure,  but 
decomposes  at  higher  temperatures,  explosively  above  100°,  the  chief,  but  not  the 
exclusive  products,  being  ammonia,  nitrogen,  and  water  : 
3NH2OH  =  NH3  +  N2  +  3H20. 

It  is  readily  soluble  in  water,  forming  an  alkaline  solution.  It  is  a  considerably 
weaker  base  than  ammonia,  from  which  it  may  be  regarded  as  being  derived  by 
the  replacement  of  a  hydrogen  atom  by  the  OH  group.  In  connexion  with  the 
weakly  basic  character  of  hydroxylamine,  it  is  an  interesting  fact  that  one  mole- 
cule of  a  monobasic  acid  can  combine  with  more  than  one  molecule  of  this  base. 
For  example,  the  hydrochloride  (NH2OH)a'HCl  is  known,  as  well  as  the  normal 
compound,  NH2OH'HC1. 

Hydroxylamine  is  a  powerful  reducing  agent.  It  immediately  precipitates 
silver  and  gold  from  solutions  of  their  salts  and  reduces  cupric  salts  in  alkaline 
solution  to  red  cuprous  oxide  : 


Under  certain  circumstances  hydroxylamine  can  act  as  an  oxidizing  agent. 
Thus  it  oxidizes  ferrous  to  ferric  hydroxide  in  an  alkaline  medium,  whereas  in 
acid  solution  ferric  salts  are  readily  reduced  to  the  ferrous  condition. 


NITROGEN-HALOGEN  COMPOUNDS  221 


COMPOUNDS  OF  NITROGEN  WITH  THE  HALOGENS 
Three  definite  compounds  of  nitrogen  with  the  halogens,  nitrogen  trichloride, 

NC1S,  nitrogen  tribromide,  NBr3,  and  an  iodide,  NH3NI3,  are  known.     They  may 

be  regarded  as  ammonia  in  which  the  hydrogen  is  partially  or  completely  dis- 

placed by  halogen  atoms.     They  are  all  extremely  explosive. 
Nitrogen  Chloride,  NC13.     Preparation  —  This  compound  is  obtained  by  the 

action  of  chlorine  upon  ammonium  chloride  : 


When  a  jar  of  chlorine  is  inverted  in  a  solution  of  ammonium  chloride  kept  at 
30°  to  40°,  oily  drops  of  the  trichloride  settle  out. 

Properties  —  Nitrogen  trichloride  is  a  dark-yellow,  oily  liquid  of  density  1.653. 
It  explodes  violently  at  the  ordinary  temperature  in  contact  with  traces  of  organic 
matter  or  of  phosphorus,  and  sometimes  spontaneously,  without  any  apparent 
cause.  It  also  explodes  on  sudden  heating,  or  on  exposure  to  sunlight.  It  boils 
at  71°  and  can  be  distilled,  but  the  operation  is  a  very  dangerous  one.  The 
vapour  strongly  attacks  the  eyes  and  the  mucous  membrane.  It  is  soluble  in 
benzene,  chloroform,  carbon  tetrachloride,  and  other  liquids,  and  these  solutions 
are  fairly  stable.  The  trichloride  is  decomposed  by  excess  of  ammonia  : 


and  there  is  therefore  no  danger  in  preparing  nitrogen  from  ammonia  and  chlorine 
(p.  202)  if  the  former  is  in  excess. 

As  would  be  anticipated  from  its  very  explosive  character,  nitrogen  chloride  is 
highly  endothermic.  In  the  formation  of  one  mol  from  its  elements,  about 
40,000  cal.  are  absorbed. 

Nitrogen  Tribromide  is  obtained  as  a  red,  extremely  explosive,  oily  liquid  by 
the  action  of  potassium  bromide  on  nitrogen  trichloride.  Its  composition  has  not 
been  conclusively  established,  but,  from  analogy  with  the  chloride,  is  generally 
assumed  to  be  represented  by  the  formula  NBr3. 

Nitrogen  Iodide,  NI3'NH8.  Preparation  —  Nitrogen  iodide  is  obtained  as  a 
brown  powder  by  the  action  of  a  concentrated  solution  of  ammonia  on  powdered 
iodine,  on  iodine  dissolved  in  alcohol,  or  on  iodine  chloride,  IC1. 

Properties  —  When  pure,  nitrogen  iodide  occurs  in  lustrous,  copper-coloured 
needles  of  density  3.5.  When  moist,  it  has  not  much  tendency  to  explode,  but 
in  the  dry  state  it  is  extremely  unstable,  even  a  touch  with  a  feather  or  the  tread 
of  a  fly  being  sufficient  to  bring  about  explosion.  The  constitution  of  nitrogen 
iodide  has  been  the  subject  of  a  great  deal  of  discussion  since  its  discovery  by 
Courtois  in  1812,  but  within  the  last  ten  years  it  has  been  definitely  established 
by  Chattaway,  Orton  and  Stevens,  and  by  Silberrad,  that  the  formula  is  as  above. 


CHAPTER  XVIII 

OXIDES  AND  OXYACIDS  OF  NITROGEN 
T  7*IVE  oxides  of  nitrogen  are  known — 

Nitrous  oxide  (hyponitrous  anhydride)       .  N2O 

Nitric  oxide         .         .         .         .        .         .  NO 

Nitrogen  trioxide  (nitrous  anhydride)          .  N2O3 

Nitrogen  peroxide NO2  or  N2O4 

Nitrogen  pentoxide  (nitric  anhydride)         .  N2O6 

Three  acids  are  known,  corresponding  with  the  three  anhydrides 
referred  to  above — 

Hyponitrous  acid H2N2O2 

Nitrous  acid HNO2 

Nitric  acid HNO3 

As  all  the  other  oxides  and  oxyacids  are  derived  directly  or  in- 
directly from  nitric  acid,  it  will  be  convenient  to  deal  first  with  this 
important  compound. 

NITRIC  ACID,  HNO3 

History — Nitric  acid  was  probably  familiar  to  the  ancient 
Egyptians.  It  was  prepared  in  the  ninth  century  by  the  alchemist 
Geber  by  heating  together  nitre  (potassium  nitrate),  alum  and  copper 
sulphate.  The  method  of  preparation  now  in  use— action  of  sulphuric 
acid  on  an  alkali  nitrate — appears  to  have  been  introduced  by  Glauber 
about  1660.  In  1784  Cavendish  obtained  it  by  passing  electric  sparks 
through  a  mixture  of  nitrogen  and  oxygen,  and  bringing  the  product 
in  contact  with  water ;  in  this*  way  the  constitution  of  the  acid  was 
established.  In  allusion  to  its  powerful  solvent  action  on  metals,  it 
was  formerly  called  aquafortis-  The  name  now  in  use  refers  to  its 
usual  mode  of  preparation  from  nitre. 

Occurrence — Nitric  acid  does  not  occur  free  in  nature,  except 
perhaps  in  traces  in  the  atmosphere,  but  in  the  form  of  its  salts,  the 


OXIDES   AND   OXYACIDS   OF   NITROGEN       223 

nitrates,  it  is  widely  distributed.  Sodium  nitrate,  NaNO3,  known  as 
Chili  saltpetre,  occurs  in  enormous  deposits  in  the  rainless  regions  of 
South  America.  Further,  nitrates  represent  the  final  oxidation  pro- 
duct of  nitrogenous  animal  and  vegetable  matter,  and  for  this  reason 
they  are  found  in  all  soils.  It  is  interesting  to  note  that  this  oxidation 
is  effected  by  atmospheric  oxygen  with  the  co-operation  of  certain 
bacteria  (p.  415).  The  occurrence  of  nitrates,  chiefly  ammonium 
nitrate,  in  the  atmosphere  has  already  been  mentioned. 

Preparation — (i)  The  preparation  of  nitric  acid  by  passing 
sparks  through  a  mixture  of  oxygen  and  nitrogen,  and  treating  the 
product  with  water,  has  already  been  referred  to.  The  reaction  may 
be  carried  out  by  means  of  the  arrangement  represented  in  Fig.  50. 
Sparks  from  a  Ruhmkorff  coil  are  passed 
through  the  air  in  the  globe  until  the  colour 
becomes  brown  owing  to  the  formation  of 
nitrogen  peroxide,  NO2 ;  on  shaking  with 
water  the  solution  will  be  found  to  have  an 
acid  reaction  : 


Nitric  acid  and  nitric  oxide  are  formed,  and 

the  reaction  is  reversible.     A  modification  of 

this  process  has  been  introduced  in  recent 

years  for  the   purpose   of  preparing    nitric    _ 

acid  on  the  large  scale  from  the  atmosphere ;  FIG.  50. 

It  is  briefly  described  on  p.  236. 

(2)  Nitric  acid  is  usually  prepared  in  the  laboratory  by  heating 
potassium  or  sodium  nitrate  with  concentrated  sulphuric  acid  in  a  glass 
retort,  and  collecting  the  acid  in  a  cooled  receiver  (Fig.  51)  : 

NaNO3  +  H2SO4->NaHSO4  +  HNO3  f  . 

Sodium  acid  sulphate,  NaHSO4,  remains  in  the  retort.  The  acid 
thus  obtained  contains  water  and  nitrogen  peroxide.  The  water  is 
almost  entirely  removed  by  adding  concentrated  sulphuric  acid  and 
redistilling,  and  the  nitrogen  peroxide  by  drawing  a  current  of  carbon 
dioxide  or  air  through  it. 

On  the  commercial  scale,  sodium  nitrate  and  sulphuric  acid  are 
mixed  in  a  cast-iron  still  in  the  proportions  required  by  the  above 
equation,  the  acid  is  distilled  off  under  reduced  pressure  in  order  to 
minimize  decomposition  (see  below),  and  is  condensed  in  a  series  of 


224     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

earthenware  pots.     If  anhydrous  sodium  nitrate  is  used,  the  distillate 
consists  of  acid  almost  free  from  water. 

Physical  Properties  —  According  to  recent  investigations,  100 
per  cent,  nitric  acid  exists  only  at  low  temperatures  in  the  form  of 
colourless  crystals,  which  melt  at  —  41.3°,  forming  a  liquid  which  is 
yellowish  in  colour  owing  to  slight  decomposition  : 


A  mixture  containing  98.6  per  cent,  of  the  acid  is  colourless  and 
moderately  stable  at  the  ordinary  temperature.  It  fumes  in  the  air, 
and  can  be  distilled  without  decomposition  under  reduced  pressure,1 
but  under  atmospheric  pressure  it  boils  at  86°  with  partial  decomposi- 


FIG.  51. 

tion  according  to  the  above  equation.  The  commercial  acid  contains 
68  per  cent,  of  real  nitric  acid,  and  is  quite  stable.  This  composition 
corresponds  with  a  constant  boiling  mixture  of  the  acid  and  water 
(cf.  hydrochloric  acid,  p.  95),  which  boils  at  120.5°  under  at- 
mospheric pressure  and  has  a  specific  gravity  of  1.414  at  15°. 
Nitric  acid  is  completely  decomposed  into  water,  nitrogen  peroxide 
and  oxygen,  according  to  the  above  equation,  when  heated  to  260°. 

Chemical  Properties — The  more  important  chemical  proper- 
ties of  nitric  acid  can  be  classified  under  three  heads — (i)  it  is  a  typical 
monobasic  acid,  combining  with  bases  tcTform  salts,  the  nitrates  ;  (2) 
on  account  of  the  readiness  with  which  it  yields  oxygen  it  is  a  power- 
ful oxidizing  agent ;  (3)  it  acts  on  certain  organic  compounds  so  that 

1  Under  reduced  pressure  the  acid  distils  at  a  much  lower  temperature. 


OXIDES  AND    OXYACIDS    OF   NITROGEN       225 

one  or  more  atoms  of  hydrogen  are  displaced  by  univalent  NO2 
groups. 

Nitrates  are  obtained  by  the  action  of  nitric  acid  on  the  oxides  or 
carbonates  of  the  metals  or  on  the  metals  themselves.  With  very 
few  exceptions,  they  are  readily  soluble  in  water.  On  heating  strongly 
they  are  all  decomposed,  the  majority  giving  off  oxygen  and  nitrogen 
peroxide  and  leaving  an  oxide  of  the  metal. 

Nitric  acid  acts  on  all  metals  except  gold,  platinum,  iridium  and 
rhodium,  nitrates  or  oxides  of  the  metal  being  formed,  whilst  the 
nitric  acid  is  reduced  to  lower  oxides  of  nitrogen.  The  vigour  of  the 
action  and  also  the  stage  to  which  the  nitric  acid  is  reduced  depend 
on  the  nature  of  the  metal,  the  concentration  of  the  acid,  the  tempera- 
ture and  other  factors.  In  contrast  to  other  acids,  hydrogen  is  very 
seldom  evolved  when  nitric  acid  acts  on  metals.1  We  may  assume 
that  it  is  used  up  in  reducing  part  of  the  nitric  acid,  being  itself 
oxidized  to  water.  The  different  stages  in  the  reduction  of  nitric 
acid  are  represented  by  nitrogen  peroxide,  NO2,  nitrogen  trioxide, 
N2O3,  or  nitrous  acid,  HNO2,  nitric  oxide,  NO,  nitrous  oxide,  N2O, 
nitrogen,  N2,  hydroxylamine,  NH2OH,  and  finally  ammonia,  NH3. 

When  copper  is  acted  on  by  nitric  acid,  the  main  reaction  is 
represented  by  the  equation 


and  the  action  on  mercury  and  on  lead  is  represented  by  correspond- 
ing equations.  With  zinc,  on  the  other  hand,  reduction  to  nitrous 
oxide  (and  under  certain  circumstances  even  to  ammonia)  takes  place  : 

4Zn+ioHNO3=4Zn  (NO3)2  +  N2O  +  5H2O. 

Altnough  gold  and  platinum  are  not  acted  on  by  nitric  acid,  they  are 
readily  dissolved  by  a  mixture  of  nitric  and  hydrochloric  acids,  the  so- 
called  aqua  regia  (p.  235). 

Nitric  acid  also  exerts  an  oxidizing  action  on  many  non-metals, 
with  formation  of  oxides  and  oxyacids.  Thus  iodine  is  oxidized  to 
iodic  acid,  HIO3  (p.  185),  phosphorus  to  phosphoric  acid,  H3PO4,  and 
sulphur  to  sulphuric  acid,  H2SO4.  The  action  of  nitric  acid  on  metals 
and  non-metals  is  further  referred  to  in  connexion  with  the  elements 
themselves. 

When  nitric  acid  acts  on  glycerine,  nitroglycerine,  the  principal 
constituent  of  dynamite,  is  obtained,  and  similarly,  by  the  action 
of  nitric  acid  on  cotton-wool,  the  highly  explosive  gun-cotton  is 

1  When  nitric  acid  acts  on  magnesium,  hydrogen  is  liberated. 
15 


226     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

obtained.  As  already  mentioned,  these  compounds  are  formed  by  the 
displacement  of  hydrogen  atoms  in  the  organic  compounds  by  NO2 
groups,  and  illustrate  the  third  main  property  of  nitric  acid.  A 
similar  kind  of  change  takes  place  when  dilute  nitric  acid  acts  on  the 
skin,  yellow  compounds  being  produced.  Concentrated  nitric  acid, 
on  the  other  hand,  causes  painful  wounds. 

Fuming  nitric  acid  is  obtained  by  distilling  concentrated  nitric  acid 
with  concentrated  sulphuric  acid,  or  by  passing  nitrogen  trioxide  into 
nitric  acid.  It  consists  of  concentrated  nitric  acid  containing  nitrogen 
peroxide  in  solution,  is  red  in  colour,  fumes  strongly  in  the  air  and 
has  powerful  oxidizing  properties. 

The  mechanism  of  the  action  of  nitric  acid  on  metals  is  not  weli 
understood.  An  important  point  in  this  connection  is  that  nitric  acid 
quite  free  from  nitrous  acid  has  very  little  action  on  copper  or  mercury. 
The  most  plausible  explanation  hitherto  advanced  is  that  hydrogen  is 
liberated  in  the  first  stage  of  the  reaction,  and  that  nitrous  acid  greatly 
increases  the  rate  of  oxidation  of  the  hydrogen  by  nitric  acid,  thus 
allowing  the  reaction  to  proceed. 

(i)  [Cu+2HNO3->Cu(NO3)2  +  H2]x3. 
(2) 
Adding  up — 

(3) 

Methods  of  Writing  Complicated  Equations— The 
above  is  an  illustration  of  writing  equations  in  stages,  a  method  which 
should  always  be  used  for  complicated  equations.  As  a  preliminary 
to  writing  the  equation,  the  stage  to  which  the  nitric  acid  is  reduced 
must  be  determined  experimentally,  and  the  formula  of  the  nitrate 
must  be  known.  In  the  above  instance  3H2  are  required  to  reduce 
2HNO3  to  nitric  oxide  and  water  ;  to  obtain  this  amount  of  hydrogen, 
equation  (i)  has  to  be  multiplied  by  3. 

A  more  complicated  case  is  that  of  the  action  of  nitric  acid  on  iodine 
to  form  iodic  acid,  the  oxide  corresponding  with  which  is  I2O5.  The 
method  of  procedure  will  be  clear  from  the  accompanying  equations 
without  further  explanation  : 

(1)  [I2+5H20->I206+5H2]x3. 

(2)  [2HNO3+3H2->4H2O  +  2NO]x5. 

(3)  [I206+H20->2HI03]X3- 

Adding — 

3l2+ioHNO3-»6HIO34-ioNO  +  2H2O. 

The  equations  just  given  probably  do  not  represent  the  mechanism 
of  the  action  of  nitric  acid  on  iodine,  but  it  should  be  understood  that 


OXIDES   AND    OXYACIDS    OF    NITROGEN       227 

we  are  dealing  with  arithmetical  methods  of  deducing  complicated 
equations  which  do  not  in  any  case  necessarily  represent  the  actual 
stages  ;  these  in  most  cases  are  quite  unknown. 

The  student  should  apply  this  method  to  other  reactions,  for  example, 
to  the  action  of  nitric  acid  on  phosphorus,  on  sulphur,  and  on  zinc. 

Another  method  of  deriving  these  equations  is  to  assume  as  the  first 
stage  the  reduction  of  the  nitric  acid  by  the  metal  or  other  element  to 
a  definite  oxide  of  nitrogen,  the  other  oxide,  if  that  of  a  metal,  then 
combines  with  more  nitric  acid  to  form  the  nitrate. 

NITROGEN  PENTOXIDE  (NITRIC  ANHYDRIDE),  N2O6 

Preparation  —  (i)  By  the  abstraction  of  the  elements  of  water 
from  nitric  acid  by  means  of  phosphorus  pentoxide  : 


Phosphorus  pentoxide  is  added  to  pure  concentrated  nitric  acid  in 
the  cold  till  a  little  remains  undissolved  ;  the  mixture  is  then  cautiously 
distilled  and  the  pentoxide  collected  in  a  cool  receiver. 

(2)  By  the  action  of  chlorine  on  silver  nitrate  : 


Dry  chlorine  is  passed  over  silver  nitrate  in  a  U-tube  kept  at  50-60°  ; 
the  pentoxide  distils  off  and  is  collected  in  a  cool  receiver. 

Properties  —  Nitrogen  pentoxide  occurs  in  colourless,  lustrous, 
rhombic  crystals,  which  melt  at  29.5°  with  partial  decomposition. 
The  reddish  liquid  begins  to  boil  at  45°  and  decomposes  rapidly, 
giving  off  brown  fumes  of  nitrogen  peroxide  : 


It  dissolves  in  water  to  form  nitric  acid,  with  evolution  of  much 
heat: 


NITROGEN  PEROXIDE,  NO2  OR  N2O4 

Preparation— (i)  By  passing  electric  sparks  through  a  mixture 
of  nitrogen  and  oxygen  : 


(2)  By  direct  combination  of  nitric  oxide  and  oxygen  : 


228    A   TEXT-BOOK   OF  INORGANIC   CHEMISTRY 
(3)  By  heating-  the  nitrate  of  a  heavy  metal,  e.g.  lead  nitrate : 
2Pb(N  O^-xzPbO  +  4NO2  +  O2. 

The  carefully  dried  salt  is  heated  in  a  glass  tube  and  the  products 
of  decomposition  passed  into  a  U-tube  immersed  in  a  mixture  of  ice 
and  salt,  the  nitrogen  peroxide  condensing  as  a  nearly  colourless 
liquid. 

Physical  Properties — At  ordinary  temperatures  nitrogen  per- 
oxide is  usually  met  with  as  a  reddish-brown  gas  which  can  readily 
be  condensed  to  a  liquid.  On  further  cooling  it  forms  colourless 
crystals,  which  melt  at  - 10°.  Even  at  its  melting-point  the  liquid 
is  pale  yellow,  and  the  colour  gradually  deepens  to  orange  as  the 
temperature  rises.  The  liquid  boils  at  26°,  changing  to  a  reddish- 
brown  vapour,  which  becomes  progressively  deeper  in  colour  as  the 
temperature  is  further  raised.  These  changes  of  colour  take  place 
in  the  reverse  order  as  the  temperature  is  lowered.  Simultaneously 
with  the  deepening  of  colour,  a  diminution  in  density  occurs  as  the 
temperature  is  raised,  as  the  following  figures  show : 

Temperature.        .     26.7°    60.2°    90.0°     121.5°     r35°     IS°° 
Density.         .        .     38.3      30.1      24.8      23.5        23.1      23.0 

Tbese  numbers  indicate  that  at  26.7°  the  molecular  weight  is  about 
76  and  at  150°  46.  The  formula  NO2  corresponds  with  a  molecular 
weight  of  46  and  the  double  formula  N2O4  with  a  molecular  weight 
of  92.  The  simplest  explanation  of  the  above  numbers  is  that  at 
150°  the  gas  consists  entirely  of  NO2  molecules  and  at  lower  tem- 
peratures of  a  mixture  of  NO2  and  N2O4  molecules.  From  the 
changes  of  colour  we  may  assume  that  the  compound  N2O4  is  colour- 
less and  NO2  deep  brown.  We  are  therefore  dealing  with  an 
equilibrium  represented  by  the  equation 

N204     ^     2N02 
I  unit  vol.      2  unit  vols. 

and  it  can  be  calculated  from  the  above  figures  that  at  26.7°  the 
mixture  contains  33  per  cent,  of  NO2  molecules  whilst  at  150°  dis- 
sociation is  practically  complete. 

On  heating  nitrogen  peroxide  to  higher  temperatures,  it  dissociates 
(p.  169)  into  nitric  oxide  and  oxygen: 


OXIDES  AND   OXYACIDS   OF  NITROGEN       229 

Under  atmospheric  pressure,  dissociation  is  practically  complete  at 
620°. 

Nitrogen  peroxide  is  a  supporter  of  combustion,  provided  the  tem- 
perature is  sufficiently  high  to  liberate  oxygen.  Thus  it  does  not 
ignite  a  glowing  splinter,  but  phosphorus  burning  in  air  continues 
to  burn  brilliantly  in  nitrogen  peroxide.  It  acts  as  a  powerful 
oxidizing  agent  towards  many  substances. 

When  water  in  moderate  excess  acts  on  nitrogen  peroxide  at  low 
temperatures  a  mixture  of  nitrous  and  nitric  acids  is  obtained  : 

2NO2  +  H2O->HNO2  +  HNO3. 
At  higher  temperatures  nitric  acid  and  nitric  oxide  are  formed  : 


With  a  cold  aqueous  solution  of  potassium  hydroxide,  a  mixture 
of  nitrite  and  nitrate  is  obtained  : 

2NO2+2KOH->KNO2  +  KNO3+H2O. 

These  facts  show  that  nitrogen  peroxide,  like  chlorine  peroxide,  is 
a  mixed  anhydride,  since  it  forms  two  acids  with  water  and  two  salts 
with  bases. 

Nitrogen  Trioxide,  N2O3—  Preparation  —  (i)  When  equal 
volumes  of  nitrogen  peroxide  and  nitric  oxide  are  passed  through 
a  tube  cooled  to  -  20°,  nitrogen  trioxide  is  obtained  as  a  blue  liquid. 

(2)  By  heating  nitric  acid  (density  1.35)  with  arsenic  trioxide,  and 
passing  the  gas  through  a  tube  which  is  kept  cool,  the  same  product 
is  obtained  : 


Properties  —  Until  quite  recently  nitrogen  trioxide  was  only 
known  in  the  liquid  form.  Below  -21°  the  liquid  is  fairly  stable,1  but 
at  higher  temperatures  it  decomposes  into  NO  and  NO2,  and  in  the 
form  of  vapour,  as  shown  by  density  measurements,  is  decomposed 
almost  completely  into  these  two  oxides  : 


Baker  (1907)  has  shown,  however,  that  if  the  liquid  trioxide  is  dried 
extremely  carefully  before  vaporizing,  it  has  the  density  38,  so  that 

1  Liquid  nitrogen  trioxide  is  deep  indigo  blue  below  -2°,  but  is  green  (and 
doubtless  partially  decomposed)  at  the  ordinary  temperature.  On  cooling  in 
liquid  air  it  forms  deep  blue  crystals,  which  melt  at  -  103°. 


230     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

no  dissociation  had  taken  place.  In  fact,  in  many  of  Baker's  experi- 
ments a  density  exceeding  38  was  observed,  which  probably  indicates 
partial  polymerization  to  N4O6  molecules. 

Nitrogen  trioxide  dissolves  in  cold  water  to  form  nitrous  acid  : 


Nitrous  Acid,  HNO2—  Preparation—  The  acid  itself  is  un- 
stable, and  has  never  been  obtained  in  the  free  condition,  but  is 
moderately  stable  in  dilute  aqueous  solution  below  o°.  It  is  obtained 
by  dissolving  nitrogen  trioxide,  N2O3,  in  cold  water  : 

N2O3+H2O-»2HN02. 

The  salts  of  nitrous  acid,  the  nitrites,  are  quite  stable.  The  nitrites 
of  the  alkali  metals,  e.g.  potassium  nitrite,  are  obtained  by  heating 
the  nitrates  alone,  or  better,  with  metallic  lead  : 

2KNO3->2KNO2  +  O2 
+  Pb-»KNO 


Properties  —  Even  in  dilute  aqueous  solution  nitrous  acid  decom- 
poses slowly  at  room  temperature,  more  rapidly  on  warming,  into 
nitric  acid,  nitric  oxide,  and  water  : 

3HN02->HN03  +  2NO  +  H20. 

For  this  reason  brown  fumes  are  given  off  when  dilute  acids  (e.g. 
dilute  sulphuric  acid)  are  added  to  solutions  of  alkali  nitrites,  the 
nitrous  acid  which  is  presumably  first  liberated  decomposing  accord- 
ing to  the  above  equation  (distinction  from  nitrates).  Nitrous  acid 
is  an  acid  of  medium  strength  (p.  187)  ;  it  is  much  weaker  than  nitric 
acid. 

When  a  strong  oxidizing  agent,  such  as  potassium  permanganate,  is 
added  to  nitrous  acid,  the  latter  is  oxidized  to  nitric  acid  : 

2HNO2+O2=2HNO3, 

and  can  therefore  act  as  a  reducing  agent.  On  the  other  hand,  nitrous 
acid  gives  up  oxygen  to  certain  easily  oxidized  substances,  such  as 
hydrogen  sulphide,  hydriodic  acid,  or  indigo  solution,  being  itself 
reduced  to  a  lower  oxide  of  nitrogen,  for  example,  nitrous  oxide  : 


According  to  the  conditions,  nitrous  acid  can  therefore  act  as  an 
oxidizing  or  as  a  reducing  agent.     In  these  respects  it  resembles 


OXIDES   AND   OXYACIDS   OF  NITROGEN       231 

hydrogen  peroxide,  which  is  also  able  to  reduce  an  acidified  solution 
of  potassium  permanganate  and  to  oxidize  indigo  and  hydriodic  acid 
(p.  141). 

NITRIC  OXIDE,  NO 

Nitric  oxide  was  first  recognized  as  a  definite  chemical  compound 
by  Priestley  (1772).  It  is  the  substance  first  formed  by  the  direct  com- 
bination of  nitrogen  and  oxygen  at  high  temperatures  or  under  the 
influence  of  the  electric  discharge. 

Preparation  —  (i)  By  the  action  of  nitric  acid  of  density  1.2  (30  to  35 
per  cent,  of  the  pure  acid)  on  copper,  the  temperature  being  kept  low  : 


The  gas  obtained  in  this  way,  although  sufficiently  pure  for  some 
purposes,  always  contains  other  oxides  of  nitrogen  in  small  amount, 
and  may  be  purified  by  passing  into  a  cold  concentrated  solution  of 
ferrous  sulphate,  which  absorbs  only  nitric  oxide  (see  below)  and  then 
heating  the  solution.  The  gas  is  collected  over  water. 

(2)  Pure  nitric  oxide  is  obtained  by  heating  nitric  acid  with  a  mixture 
of  ferrous  sulphate,  FeSO4,  and  dilute  sulphuric  acid,  or  ferrous  chlo- 
ride and  hydrochloric  acid  : 


[2FeSO4  +  H2SO4  +  O->Fe2(SO4)3  +  H2O]  x  3. 
Adding  — 


Fe2(SO4)3  is  ferric  sulphate  (p.  562). 

Physical  Properties—  Nitric  oxide  is  a  colourless  gas.  Its 
density  referred  to  hydrogen  is  15,  corresponding  with  the  formula 
NO  ;  unlike  nitrogen  peroxide,  it  shows  no  tendency  to  polymeriza- 
tion, even  at  -  70°.  It  can  be  obtained  as  a  colourless  liquid,  which 
boils  at  -153.6°;  the  critical  temperature  is  -93.  5°  and  the  critical 
pressure  71.2  atmospheres  (Olszewski).  Solid  nitric  oxide  occurs  in 
colourless  crystals,  which  melt  at  —  167°. 

Nitric  oxide  is  very  slightly  soluble  in  water.  At  o°  i  volume  of 
water  absorbs  0.074  c.c.,  at  10°  0.057  c.c.,  and  at  20°  0.647  c.c.  of 
the  gas. 

As  nitric  oxide  is  partially  decomposed  into  its  elements  at  high 
temperatures  in  the  presence  of  catalytic  agents,  the  reaction 


is  reversible.      Nernst  has  shown  that  at    1760°  C.  the  equilibrium 


232     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

mixture  contains  0.64  per  cent.,  at  2210°  2.05  per  cent.,  and  at  3000° 
about  5  per  cent,  of  nitric  oxide.  The  displacement  of  the  equilibrium 
in  favour  of  nitric  oxide  with  increasing  temperature  is  in  accordance 
with  its  endothermic  character,  the  thermochemical  equation  being 

N2+O2=2NO  -2x21,  500  cal. 

It  follows  that  at  ordinary  temperatures  nitric  oxide  is  unstable,  its 
apparent  stability  being  due  to  the  fact  that  the  reaction  2NO->N2  +  O2 
is  exceedingly  slow  under  these  conditions. 

Chemical  Properties  —  Owing  to  the  excessive  slowness  with 
which  it  splits  up  into  its  elements,  nitric  oxide  is  the  most  stable  of 
all  the  oxides  of  nitrogen.  Even  at  a  red  heat  it  is  practically 
unaffected.  For  the  same  reason  it  is  not  a  supporter  of  combustion 
in  the  ordinary  sense,  as  a  lighted  candle  or  phosphorus  burning 
feebly  are  extinguished  when  plunged  into  it.  If,  however,  the  tem- 
perature is  sufficiently  high  to  decompose  it,  it  supports  combustion  ; 
thus  burning  magnesium  or  brightly  burning  phosphorus  continue  to 
burn  in  it. 

As  already  mentioned,  nitric  oxide  combines  with  oxygen  to  form 
brown  fumes  of  nitrogen  peroxide  : 


The  combination  is  practically  complete  at  low  temperatures,  but 
above  1  50°  dissociation  of  the  peroxide  becomes  appreciable. 

Nitric  oxide  is  readily  absorbed  by  aqueous  solutions  of  ferrous 
salts,  e.g.  ferrous  sulphate,  forming  a  dark-coloured  compound  in 
solution  which  is  readily  decomposed  on  heating  (p.  231).  The  consti- 
tution of  this  compound  has  not  been  satisfactorily  established,  but  is 
probably  FeSO4'NO  (Manchot  and  Zechentmayer,  1907).  On  the 
formation  of  this  compound  is  based  the  so-called  "  brown  ring  "  test 
for  nitrates.  To  a  solution  of  a  nitrate  a  crystal  of  ferrous  sulphate 
is  added,  and  after  shaking  and  pouring  concentrated  sulphuric  acid 
down  the  inside  of  the  tube,  a  brown  ring  of  the  above  compound 
forms  at  the  junction  of  acid  and  solution. 

Composition  —  Nitric  oxide  is  completely  decomposed  on  heat- 
ing with  metallic  sodium  or  iron,  and  the  residual  nitrogen  has  just 
half  the  volume  of  the  original  gas  : 


2  VOls.  I  Vol. 


A  molecule  of  nitric  oxide,  therefore,  contains  one  atom  of  nitrogen, 


OXIDES   AND   OXYACIDS   OF   NITROGEN      233 

weight  14,  and  as  the  density  of  nitric  oxide  is  15,  and  therefore  its 
molecular  weight  30,  the  molecule  contains  30-14  =  16  parts  by 
weight,  or  one  atom  of  oxygen,  and  its  formula  is  therefore  NO. 


NITROUS  OXIDE  (NITROUS  ANHYDRIDE,  LAUGHING  GAS),  N2O 

This  gas  was  first  obtained  by  Priestley  (1772)  by  the  action  of  iron 
filings  in  contact  with  water  on  nitric  oxide ;  its  composition  and 
more  important  properties  were  established  by  Davy. 

Preparation — On  cautiously  heating  ammonium  nitrate  (the 
salt,  not  a  solution)  it  splits  up  directly  into  water  and  nitrous  oxide  : 

NH4NOr*N1O+2H,O. 

Instead  of  ammonium  nitrate  a  mixture  of  ammonium  sulphate  and 
sodium  nitrate  in  equivalent  proportions  may  be  used  ;  by  double 
decomposition  ammonium  nitrate  is  formed,  which  then  decomposes 
as  above.  Ammonium  chloride  should  not  be  used  instead  of  ammo- 
nium sulphate,  as  a  very  impure  gas  is  then  obtained. 

When  required  pure,  e.g.  for  anaesthetic  purposes,  the  gas  is  passed 
through  a  solution  of  ferrous  sulphate  to  remove  nitric  oxide,  through 
alkali  to  remove  traces  of  chlorine,  and  is  collected  over  hot  water 
or  over  mercury. 

Nitric  acid  can  be  reduced  to  nitrous  oxide  by  the  action  of  metals 
such  as  zinc  or  copper  under  certain  conditions,  but  a  mixture  of  gases 
is  always  obtained  in  these  reactions  (p.  225). 

Physical  Properties — Nitrous  oxide  is  a  colourless  gas  with 
an  agreeable  sweetish  odour  and  taste.  Its  density  is  22.  It  can 
readily  be  condensed  to  a  colourless  liquid,  which  boils  at  -89.5°; 
its  critical  temperature  is  35.4°  and  critical  pressure  75  atmospheres. 
The  liquid  is  obtainable  commercially,  compressed  in  steel  cylinders. 
Solid  nitrous  oxide  occurs  in  colourless  crystals,  which  melt  at 
- 102.3. 

It  was  first  observed  by  Davy  that  when  nitrous  oxide  is  inhaled 
in  small  amount,  it  gives  rise  to  excitement  often  accompanied  by 
hysterical  laughter ;  hence  the  name  laughing  gas.  On  continued 
inhalation  it  induces  insensibility,  and  is  therefore  used  as  an  anaes- 
thetic, especially  in  dental  operations. 

Nitrous  oxide  is  fairly  soluble  in  water.  One  c.c.  of  water  at  5° 
absorbs  1.048  c.c.,  at  15°  0.7377  c.c.,  at  25°  0.5443  c.c.  of  the  gas. 

Nitrous  oxide  is  decomposed  into  its  elements  at  a  much  lower 
temperature  than  is  nitric  oxide ;  at  700°  the  speed  of  the  reaction 


234     A  TEXT  BOOK   OF   INORGANIC   CHEMISTRY 

can  readily  be  measured.  Like  nitric  oxide,  nitrous  oxide  is  an 
endothermic  compound,  about  20,000  calories  being  absorbed  in 
the  formation  of  30  grams  of  the  gas  from  its  elements,  and  its 
apparent  stability  at  ordinary  temperatures  is  due  to  its  slow  rate 
of  decomposition.  The  direct  formation  of  nitrous  oxide  from  its 
elements  has  not  hitherto  been  observed. 

Chemical  Properties  —  Owing  to  its  relatively  slight  stability, 
nitrous  oxide  is  a  supporter  of  combustion,  a  glowing  splinter  plunged 
into  it  bursts  into  flame,  and  brightly  burning  sulphur  and  phos- 
phorus (though  not  feebly  burning  sulphur)  continue  to  burn  in 
it.  It  can  at  once  be  distinguished  from  oxygen,  however,  by  the 
fact  that  no  red  fumes  are  produced  when  it  is  mixed  with  nitric 
oxide.  It  does  not  combine  directly  with  oxygen  even  at  a  red  heat. 

Composition  —  When  a  confined  volume  of  nitrous  oxide  is 
heated  with  metallic  sodium,  sodium  oxide  and  nitrogen  are  formed, 
and  the  volume  of  the  nitrogen  is  equal  to  that  of  the  nitrous  oxide 
taken.  It  follows  that  the  molecule  of  nitrous  oxide  contains  a 
molecule  or  two  atoms  of  nitrogen,  and  its  formula  must  be  N2OX, 
where  x  is  a  whole  number.  As,  however,  its  density  is  22  and  mole- 
cular weight  44,  it  contains  44-  28=  16  parts  or  i  atom  of  oxygen,  and 
its  formula  is  therefore  N2O. 

The  same  conclusion  is  also  reached  by  exploding  nitrous  oxide 
with  an  equal  volume  of  hydrogen.  Water  is  formed,  and  a  volume 
of  nitrogen  equal  to  that  of  the  original  gas  liberated: 

N20     +     H2     =     N2     +     H20 
(l  vol.)       (l  vol.)     (i  vol.)      (liquid). 

Hyponitrous  Acid—  Preparation—  (i)  The  alkali  salts  of  this  acid,  the  hypo- 
nitrites,  are  formed  when  nascent  hydrogen  (obtained  by  the  action  of  water  on 
sodium  amalgam)  acts  on  an  alkali  nitrate,  or,  better,  oh  a  nitrite  : 


The  acid  itself  cannot  be  obtained  by  addition  of  sulphuric  acid  to  a  hyponi- 
trite  in  aqueous  solution,  as  it  immediately  decomposes  into  nitrous  oxide  and 
water:  2KNO  +  H2SO4->K2SO4  +  N2O+H2O  ;  but  is  obtained  by  adding  silver 
hyponitrite  to  a  solution  of  hydrogen  chloride  in  ether,  filtering  and  evaporating 
off  the  ether  in  a  desiccator  : 


(2)  Hyponitrous  acid  is  also  obtained,  though  not  in  very  good  yield,  by  the 
action  of  nitrous  acid  on  hydroxylamine  : 

HO  -NH2+  ONOH-^HO  -N:N  -OH  +  H20, 
a  reaction  which  would  appear  to  establish  its  constitution  as  shown. 


OXIDES   AND   OXYACIDS   OF   NITROGEN       235 

Properties  —  Hyponitrous  acid  occurs  in  colourless  crystalline  leaflets,  and 
is  highly  explosive.  It  is  readily  soluble  in  water,  forming  a  moderately 
stable  solution  ;  but  decomposes  into  nitrous  oxide  and  water  on  warming. 
Freezing-point  determinations  in  aqueous  solution  prove  that  the  acid  has  the 
double  formula,  H2N2O2.  Nitrous  oxide  does  not  form  hyponitrous  acid  with 
water,  so  that  the  reaction  H2N2(V->N2O  +  H2O  is  not  reversible. 

Nitramide,  NH2NO2,  an  isomer  of  hyponitrous  acid,  is  obtained  by  the  action 
of  cold  sulphuric  acid  on  potassium  nitrocarbamate  : 


It  forms  colourless  crystals,  which  melt  at  72-75°  and  are  readily  soluble  in  water  ; 
the  solution  is  strongly  acid.  Under  the  influence  of  catalytic  agents  it  splits 
up  into  nitrous  oxide  and  water.  Its  graphic  formula  has  not  been  definitely 
settled. 


COMPOUNDS  OF  OXYGEN,  NITROGEN  AND  THE  HALOGENS 

Nitrosyl  Chloride,  NOCI  —  Preparation  —  (i)  By  direct  combination  of  nitric 
oxide  and  chlcrine  : 


(2)  A  mixture  of  I  part  nitric  acid  (density  1.2)  and  and  3  parts  hydrochloric 
acid  (density  1.12)  is  known  as  aqua  regia,  because  it  possesses  the  property 
of  dissolving  the  so-called  '  '  noble  "  metals  gold  and  platinum.  It  owes  its 
activity  to  free  chlorine,  but  also  contains  nitrosyl  chloride  : 


From  the  mixture  nitrosyl  chloride  can  be  separated  by  fractional  distillation. 

Properties  —  Nitrosyl  chloride  at  ordinary  temperature  is  an  orange-yellow 
gas,  which  is  readily  condensed  to  a  reddish  liquid  which  boils  at  —5.6°.  It 
begins  to  decompose  into  its  components  only  when  heated  to  700°.  It  is 
readily  hydrolyzed  by  water,  with  formation  of  nitrous  and  hydrochloric  acids  : 

NOC1+  HOH-»HN02+  HC1. 

Nitrosyl  Bromide,  NOBr,  formed  by  direct  combination  of  nitric  oxide  and 
bromine,  is  a  brownish  -black  liquid  at  low  temperatures.  It  boils  at  -3°,  and 
at  room  temperature  the  vapour  is  already  partially  decomposed  into  nitric  oxide 
and  bromine. 

Nitrosyl  Fluoride,  NOF,  is  formed  when  nitrosyl  chloride  is  passed  over  silver 
fluoride  heated  to  200-250°  in  a  platinum  tube  : 

NOCI  +  AgF_>NOF  +  AgCl. 

Nitrosyl  fluoride  is  a  colourless  gas,  which  can  be  condensed  to  a  colourless 
liquid  boiling  at  -56°.     It  is  immediately  hydrolyzed  by  water  with  formation  of 
nitrous  and  hydrofluoric  acids. 
A  Nitryl  Fluoride,  NO2F,  has  also  been  described. 


236     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Utilization  of  Atmospheric  Nitrogen1  —  As  nitrogen  is 
one  of  the  elements  essential  for  the  growth  of  plants  and  is  con- 
tinually being  removed  in  the  course  of  agricultural  operations,  it 
becomes  a  matter  of  the  utmost  importance  to  discover  means  of 
keeping  up  the  supply.  Ammonium  salts  and  nitrates,  especially 
sodium  nitrate  (Chili  saltpetre),  have  long  been  and  still  are  in  use 
as  nitrogenous  manures,  but  quite  recently  methods  have  been 
devised  for  utilizing  the  almost  unlimited  stores  of  nitrogen  in  the 
atmosphere.  The  most  promising  of  these  methods  is  based  upon 
the  combination  of  nitrogen  and  oxygen  in  the  electric  arc  to  nitric 
oxide.  According  to  the  Birkeland-Eyde  method,  used  in  Norway, 
the  electric  arc  is  drawn  out  by  means  of  powerful  electro-magnets 
into  a  disk  of  flame  several  feet  in  diameter.  Air  is  passed  through  the 
arc  thus  formed,  and  at  the  high  temperature  (about  3000°)  a  certain 
proportion  of  nitric  oxide,  about  I  per  cent  of  the  gas  leaving  the 
tube,  is  formed  (p.  231).  The  gases  are  cooled  and  passed  into  the 
oxidation  chambers,  in  which  the  nitric  oxide  is  completely  converted 
to  peroxide  (p.  232),  and  are  then  passed  up  towers  filled  with 
quartz  or  coke,  over  which  water  trickles,  the  nitrogen  peroxide 
being  thereby  almost  completely  transformed  to  nitric  acid  : 


The  nitric  oxide  combines  with  more  oxygen  and  is  again  passed  into 
the  absorption  tower,  till  finally  absorption  is  practically  complete. 

The  nitric  acid  is  treated  with  calcium  carbonate  or  hydroxide, 
forming  basic  calcium  nitrate,  which  is  sold  as  a  manure,  and  is  also 
used  in  the  manufacture  of  nitric  acid. 

Another  method  of  utilizing  atmospheric  nitrogen  depends  upon 
the  fact  that  when  impure  calcium  carbide  is  heated  in  nitrogen  at 
1000°  the  latter  is  absorbed  with  formation  of  a  mixture  of  calcium 
cyanamide  and  carbon  : 


The  reaction  is  exothermic  in  the  direction  of  the  upper  arrow  and 
is  reversible.  The  crude  mixture  of  calcium  cyanamide  and  carbon 
is  known  as  nitrolim.  Its  importance  depends  on  the  fact  that  the 
cyanamide  reacts  with  water  to  give  calcium  carbonate  and  ammonia  : 


1  An  excellent  account  of  the  utilization  of  atmospheric  nitrogen,  with  illustra- 
tions, is  given  by  Crossley,  Pharm.  Journal,  1910,  84,  329. 


OXIDES   AND   OXYACIDS   OF   NITROGEN      237 


and  hence  it  may  be  used  directly  as   a  manure.     In  the  soil  the 
decomposition  of  nitrolim  is  doubtless  more  complicated. 

Within  the  last  few  years  a  number  of  factories  have  been 
established  for  the  manufacture  of  nitrolim,  the  largest  being  in 
Norway,  where  water  power  is  cheap.  The  carbide  is  made  by 
heating  calcium  oxide  and  anthracite  coal  in  the  electric  furnace 
(p.  331),  the  nitrogen  by  the  liquefaction  and  subsequent  fractionation 
of  air  according  to  the  Linde  process. 

Nitrolim  is  now  used  for  the  commercial  production  of  cyanides ; 
for  this  purpose,  as  it  contains  the  requisite  substances,  it  is  only 
necessary  to  heat  it  with  a  flux J  : 

CaNCN  +  C->Ca(CN)2. 

The  Nitrogen  Cycle  in  Nature— The  nitrogen  in  nature,  like  the  carbon 
(p.  351),  is  passing  through  a  continuous  cycle  of  changes,  in  the  course  of  which 
it  exists  in  combination  in  a  number  of  forms  from  the  highly  reduced  condition 
in  ammonia  to  the  highly  oxidized  condition  in  nitrates.  All  living  material  and 
the  waste  products  of  animals  contain  combined  nitrogen  in  considerable  amount. 
When  animal  and  vegetable  matter  decays  it  undergoes  simplification  under  the 
influence  of  bacteria ;  part  of  the  nitrogen  is  given  off  as  nitrogen  gas,  part  is 
changed  into  ammonia,  and  ultimately,  by  the  agency  of  nitrifying  bacteria,  into 
nitrites,  and  finally  nitrates.  The  combined  nitrogen  essential  for  the  growth  of 
plants  and  of  animals,  which  are  dependent  on  plants  for  their  nitrogen,  is  obtained 
from  the  products  of  decay  of  animal  and  vegetable  matter,  from  manures  supplied 
to  the  soil,  from  ammonia  and  nitrates  formed  by  electric  discharge  in  the  atmos- 
phere (p.  205),  and  from  the  atmosphere  in  the  case  of  leguminous  plants  (p.  206). 

The  following  diagram,  due  to  von  Braun,  summarises  in  a  very  striking  way 
the  nitrogen  cycle  in  nature. 


fi 

z  -a 

tt 

2!" 
>  Z 

r  3 

1 

.£ 

1" 

ON!A       | 

NITRIFYING  BACTERIA    V 

1  A  flux  is  a  readily  fusible  material  which  is  added  to  infusible  substances  in 
order  to  bring  them  more  intimately  into  contact  at  high  temperatures,  chemical 
combination  being  thereby  facilitated. 


CHAPTER  XIX 

PHOSPHORUS 

Symbol,  P.     Atomic  weight  =31.0.     Molecular  weight  =  124. 


Relations  —  The  compounds  of  phosphorus  are  in 
_>4  many  cases  of  the  same  type  as  those  of  nitrogen,  due  to  the 
fact  that  the  former  element,  like  the  latter,  is  mainly  trivalent  and 
pentavalent.  The  best-known  hydrogen  compound  of  phosphorus  is 
PH3,  corresponding  with  NH3,  and  phosphorus  forms  two  principal 
oxides,  P2O3  (or  P4Og)  and  P2O6,  which,  like  the  corresponding  oxides 
of  nitrogen,  are  acidic.  From  the  first-mentioned  oxide  is  derived 
phosphorous  acid,  H3PO3,  and  the  corresponding  salts,  the  phos- 
phites ;  from  P2O6  is  derived  phosphoric  acid,  H3PO4,  and  the 
phosphates.  The  formulae  just  given  show  that  the  oxyacids  and 
oxysalts  of  phosphorus  are  not  quite  of  the  same  type  as  nitrous  and 
nitric  acids  and  the  corresponding  salts. 

History  —  Phosphorus  was  discovered  in  1669  by  Brand  of 
Hamburg,  who  obtained  it  by  evaporating  urine  (which  contains 
phosphates)  to  dryness  and  distilling  the  residue.  The  process  was 
kept  secret  for  some  years,  but  was  independently  discovered  by 
Kunkel  and  by  Boyle  (1680)  ;  the  latter  obtaining  phosphorus  by 
distilling  urine  with  sand.  The  method  now  in  use,  which  depends 
upon  the  use  of  bone-ash  (impure  calcium  phosphate)  was  discovered 
by  Scheele  (1770).  The  name  "phosphorus"  was  at  that  time 
used  to  designate  any  substance  which  glowed  in  the  dark,  and  the 
element  now  under  consideration  was  known  as  Brand's  phosphorus, 
ox  phosphorus  mirabilis. 

Occurrence  —  Phosphorus  is  never  found  free  in  nature,  but  in 
the  form  of  compounds,  chiefly  phosphates,  is  very  widely  distributed. 
Calcium  phosphate,  C^^PQ^^phosphorife,  occurs  in  large  deposits, 
and  apatite,  3Ca3(PO4)2,CaCl2  (or  CaF2)  is  found  in  volcanic  rocks. 
Other  mineral  phosphates  are  ivavellite,  4A1PO4,2A1(OH)3,9H2O,  and 
irivianite,  Fe3(PO4)2,8H2O. 

The  phosphates  resulting  from   the  disintegration  of  rocks  are 

washed  into  the  soil,  and  form  an  essential  constituent  for  the  growth 

338 


PHOSPHORUS 


239 


of  plants.     The  rigidity  of  bones  is  due  to  the  presence  of  calcium 
phosphate,  and  bone  ash  consists  almost  entirely  of  this  compound. 

Preparation  —  (i)  On  the  commercial  scale,  bone  ash,  or 
naturally  occurring  calcium  phosphate,  is  treated  with  sufficient 
sulphuric  acid  (density  1.5  to  1.6)  to  convert  it  completely  into  phos- 
phoric acid  : 


The  phosphoric  acid  is  filtered  off  from  the  insoluble  calcium  sul- 
phate, evaporated  to  a  syrupy  consistency,  mixed  with  about  one- 
quarter  of  its  weight  of  coal  or  coke,  and  heated  to  dryness  in  an 
iron  vessel.  In  this  process 
part  of  the  water  is  driven 
off,  and  metaphosphoric  acid, 
HPO3,  remains  : 

H3PO4->HPO3  +  H2O. 

The  dried  mass  is  then  put 
in  earthenware  retorts,  con- 
nected to  pipes  dipping  under 
water,  and  raised  to  a  white 
heat,  when  the  following  re- 
action takes  place  : 

4HPO3+i2C 

=  2H2+I2CO  +  P4. 

The  phosphorus  condenses 
under  water,  the  hydrogen 
and  carbon  monoxide  pass- 
ing off. 


FIG  52. 


(2)  The  preparation  is  now  effected  more  simply  by  means  of  the 
electric  furnace  (Fig.  52).  A  mixture  of  calcium  phosphate,  sand 
(silicon  dioxide,  p.  367)  and  coal  is  exposed  to  the  high  temperature 
of  the  electric  arc  passing  between  the  electrodes  as  shown,  when  a 
reaction  takes  place  represented  by  the  equation 

Ca3(PO4)2  +  3SiO 


The  sand  combines  with  the  calcium  to  form  calcium  silicate,  CaSiO3, 
a  readily  fusible  compound,  which  can  be  drawn  off  at  intervals 
through  the  opening  b.  The  phosphorus  passes  off  through  the  pipe 
«,  and  is  condensed  under  water  in  the  usual  way.  As  the  reaction 
proceeds,  fresh  charges  of  the  reacting  substances  can  be  added,  and 


24o     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  process  is  thus  made  continuous.  It  may  be  pointed  out  that 
the  process  is  not  one  of  electrolysis  ;  the  function  of  the  electric  arc 
is  simply  to  give  a  very  high  temperature. 

The  crude  phosphorus  obtained  by  the  above  methods  is  very 
impure,  being  often  red  or  almost  black  in  colour.  By  melting  and 
stirring  under  water  containing  potassium  bichromate  and  sulphuric 
acid,  part  of  the  impurities  are  oxidized  away  and  others  rise  to  the 
surface  of  the  water.  Mechanical  impurities  are  best  removed  by 
melting  under  water  and  forcing  through  chamois  leather.  The 
purified  phosphorus  is  then  cast  into  sticks.  In  order  to  prevent 
oxidation  it  is  kept  under  water. 

Physical  Properties  —  The  modification  of  phosphorus 
obtained  as  above  is  a  transparent,  almost  colourless  (slightly 
yellowish)  wax-like  solid,  of  density  1.82  at  20°.  At  o°  it  is  brittle, 
at  room  temperature  it  can  be  cut  like  wax,  under  water  it  melts 
about  44°,  forming  a  liquid  which  can  be  very  considerably  super- 
cooled without  solidifying  (p.  70).  In  absence  of  air  it  boils  at  290°. 
Even  at  room  temperature  it  is  distinctly  volatile,  as  is  evident  from 
the  characteristic  smell. 

Up  to  1000°  the  vapour  density  of  phosphorus  is  62,  correspond- 
ing with  the  formula  P4  ;  in  the  neighbourhood  of  1 500°,  however,  it 
is  considerably  less,  indicating  partial  dissociation,  probably  into  P2 
molecules. 

Yellow  phosphorus  is  insoluble  in  water,  slightly  soluble  in  alcohol, 
fairly  soluble  in  ether  and  benzene,  readily  soluble  in  carbon  disul- 
phide.  On  evaporating  the  solution  in  the  latter  solvent,  out  of 
contact  with  air,  phosphorus  is  obtained  in  well-formed  crystals 
(rhombic  dodecahedra).  As  shown  both  by  the  boiling-point  and 
freezing-point  methods  (p.  196),  phosphorus  in  solution  has  also  the 
formula  P4. 

Chemical  Properties — When  exposed  to  the  air,  yellow 
phosphorus  is  slowly  oxidized,  chiefly  to  phosphorous  oxide,  P4O6, 
which  forms  white  fumes,  the  process  being  accompanied  by  a 
garlic  odour  and  a  slight  luminosity  visible  in  the  dark.  During  the 
slow  oxidation  of  phosphorus,  ozone  (p.  134),  hydrogen  peroxide,  and 
other  products  are  formed.  Dry  phosphorus  does  not  combine  with 
perfectly  dry  oxygen,  and,  what  is  still  more  remarkable,  there  is  a 
definite  pressure  (which  depends  on  the  temperature)  above  which 
combination  with  moist  oxygen  does  not  occur,  although  it  takes 
place  readily  when  the  oxygen  pressure  is  reduced. 

Phosphorus  catches  fire  in  the  air  when  heated  to  about  35",  but 


PHOSPHORUS  241 

in  the  finely-divided  form  it  catches  fire  at  room  temperature.  This 
can  be  strikingly  shown  by  soaking  a  piece  of  filter-paper  with  a 
solution  of  phosphorus  in  carbon  disulphide  and  exposing  to  the  air  ; 
when  the  solvent  has  evaporated  the  phosphorus  bursts  into  flame. 

Yellow  phosphorus  is  a  very  poisonous  substance,  o.i  of  a  gram 
being  usually  fatal  for  an  adult.  Exposure  to  the  vapour  of  phos- 
phorus ultimately  causes  necrosis  of  the  jawbones  and  teeth  and 
other  injurious  effects. 

Red  Phosphorus.  Preparation  —  This  element  also  exists 
in  a  second  modification,  which  can  be  obtained  by  heating  yellow 
phosphorus  in  absence  of  air  at  240-250°.  On  the  commercial  scale, 
yellow  phosphorus  is  heated  in  a  closed  iron  vessel  for  some  time  at 
240°  and  then  at  a  rather  higher  temperature  in  order  to  complete 
the  change.  The  product  is  then  ground  under  water,  boiled  with 
sodium  hydroxide  to  remove  traces  of  yellow  phosphorus,  then  washed 
with  water  and  dried. 

The  change  of  yellow  to  red  phosphorus  can  also  be  brought  about 
in  solution  ;  for  example,  by  boiling  a  solution  of  yellow  phosphorus 
in  phosphorus  tribromide.  The  change  is  also  produced  under  the 
influence  of  light,  and  the  reddish  appearance  which  sticks  of 
ordinary  phosphorus  sometimes  present  is  doubtless  due  to  a  coat- 
ing of  red  phosphorus. 

The  reaction  yellow  phosphorus'->red  phosphorus  is  strongly  exo- 
thermic, about  27,300  cal.  being  given  out  in  the  transformation  of  31 
grams,  so  that  red  phosphorus  is  the  stable  form  at  high  temperatures, 
and,  judging  from  the  relative  solubilities,  also  at  ordinary  temperatures.1 

Physical  Properties — Red  or  "amorphous"  phosphorus  is  a  dark 
to  violet-red  powder  of  density  2.15.  It  is  not  luminous,  has  no 
taste  or  smell,  is  insoluble  in  carbon  disulphide,  is  not  poisonous, 
and  does  not  ignite  in  the  air  below  200° — in  all  these  respects  pre- 
senting the  most  striking  contrast  to  yellow  phosphorus.  At  fairly 
high  temperatures  red  phosphorus  is  vaporized,  and  the  vapour 
condenses  as  yellow  phosphorus.  In  a  vacuum  at  100°,  however, 
red  phosphorus  can  be  sublimed  unchanged. 

Red  phosphorus  is  sometimes  termed  amorphous  phosphorus,  but 
it  consists,  at  least  in  part,  of  minute  crystals.  It  is  probably  not  a 
uniform  substance,  and  at  present  its  exact  nature  is  not  understood. 
Cohen  (1910)  regards  red  phosphorus  as  a  "solid  solution"  (p.  198) 
of  yellow  phosphorus  in  Hittorf's  phosphorus. 

1  The  solubility  of  the  unstable  modification,  like  its  vapour  pressure  (p.  69), 
is  always  greater  than  that  of  the  stable  modification. 
16 


242     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Although  chemically  less  active  than  the  yellow  modification,  red 
phosphorus  is  readily  attacked  by  the  halogens  and  by  nitric  acid  (see 
below)  ;  it  is  only  slowly  attacked  by  boiling  alkali  hydroxide. 

Hittorf's  Phosphorus—  Hittorf  obtained  a  definite  modifica- 
tion of  this  element  by  heating  red  phosphorus  with  fused  metallic  lead 
in  a  sealed  tube.  On  cooling,  the  phosphorus  separates  in  yellowish- 
red  transparent  crystals  of  density  2.34.  This  modification  is  known 
as  metallic  phosphorus  or  Hittorf's  phosphorus. 

Matches  —  The  chief  commercial  use  of  phosphorus  is  in  the 
preparation  of  matches.  Ordinary  matches  are  made  by  dipping  the 
heads  first  into  melted  paraffin  (or  fused  sulphur)  and  then  into 
a  paste  consisting  of  yellow  phosphorus,  an  oxidizing  agent  (potassium 
chlorate,  red  lead,  or  manganese  dioxide),  and  glue.  When  rubbed 
on  a  rough  surface  (e.g.  sand-paper)  the  phosphorus  catches  fire,  and 
by  the  agency  of  the  sulphur  or  paraffin  the  wood  then  becomes 
ignited.  These  matches  are  very  poisonous  on  account  of  the 
presence  of  phosphorus,  and  for  the  same  reason  are  very  injurious  to 
the  health  of  the  workmen.  These  disadvantages  are  not  present  in 
the  so-called  "  safety  "  matches,  which  are  now  very  largely  used.  The 
heads  are  composed  of  a  mixture  of  potassium  chlorate  or  dichro- 
mate,  antimony  trisulphide,  a  little  powdered  glass  to  give  greater 
friction,  and  glue,  and  they  are  ignited  by  rubbing  on  a  prepared 
surface  containing  red  phosphorus,  antimony  trisulphide,  and  glue. 
At  the  high  temperature  produced  by  friction  the  phosphorus  burns 
in  the  oxygen  of  the  oxidizing  agents,  thus  setting  the  match  on  fire. 

COMPOUNDS  OF  PHOSPHORUS  AND  HYDROGEN 

At  least  three  compounds  of  phosphorus  and  hydrogen  are  known, 
namely,  gaseous  hydrogen  phosphide,  PH3,  liquid  hydrogen  phos- 
phide, P2H4,  and  a  solid  phosphide  (P4H2)»  ,  probably  Pi2H6.  None 
of  them  can  be  obtained  by  direct  combination  of  the  elements. 

Gaseous  Hydrogen  Phosphide  or  Phosphine,  PH3. 
Preparation  —  (i)  By  boiling  phosphorus  with  sodium  hydroxide  in 
absence  of  air: 


The  second  product  of  the  reaction,  NaH2PO2,  is  sodium  hypo- 
phosphite.  The  arrangement  of  the  apparatus  is  shown  in  Fig.  53. 
Phosphorus  and  a  concentrated  solution  of  sodium  hydroxide  are 
placed  in  the  flask,  from  which  the  air  is  then  removed  by  a  stream 


PHOSPHORUS 


243 


of  coal-gas  passed  through  the  tube.  Heat  is  then  applied,  and  as 
the  gas  escapes  into  the  air  each  bubble  immediately  catches  fire, 
forming  vortex  rings  of  metaphosphoric  acid.  As  a  matter  of  fact, 
the  spontaneous  inflammability  is  due  to  the  presence  of  traces  of 
P2H4)  and  if  the  latter  compound  is  removed  by  passing  through 


FIG.  53. 

alcohol  or  hydrochloric  acid,  the  gas  no  longer  catches  fire  at  the 
ordinary  temperature. 

(2)  By  the  action  of  water  or  of  dilute  hydrochloric  acid  on  calcium 
phosphide  : 

Ca3P2  +  6H2O->2PH3+3Ca(OH)2. 

Owing  to  secondary  reactions  traces  of  the  other  hydrides  are  also 
obtained,  and  the  gas  is  spontaneously  inflammable. 

(3)  Pure  phosphine  is  obtained  by  the  action  of  water  or  potassium 
hydroxide  on  phosphonium  iodide,  PH4I  (see  below): 


244     A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 

Physical  Properties — Phosphine  is  a  colourless  gas  with  a  dis- 
agreeable odour,  recalling  that  of  decaying  fish.  The  liquefied 
gas  boils  at  -86°.  It  is  very  slightly  soluble  in  water,  more  soluble 
in  alcohol.  It  readily  decomposes  into  its  elements  on  heating. 

Chemical  Properties — When  heated  to  about  100°  in  the  air 
phosphine  burns  to  water  and  metaphosphoric  acid : 

PH3  +  2O2  =  HPO3+H2O. 

It  does  not  combine  with  oxygen  at  atmospheric  pressure,  but  com- 
bination takes  place  explosively  when  the  pressure  is  reduced.  In 
this  respect  it  resembles  phosphorus,  which  combines  with  oxygen 
only  below  a  certain  limit  of  pressure,  which  in  both  cases  depends 
on  the  temperature  and  the  proportion  of  moisture  present. 

The  aqueous  solution  of  phosphine  does  not  affect  litmus,  so  that  it 
is  much  less  basic  than  ammonia.  It  has,  however,  a  slightly  basic 
character,  as  shown  by  its  capacity  to  combine  directly  with  the 
halogen  acids,  HX,  forming  so-called  phosphonium  compounds, 
PH4X  (analogous  to  ammonium  halogen  compounds,  NH4X),  which 
are  dealt  with  in  the  next  section.  Unlike  ammonia,  phosphine 
does  not  form  salts  with  oxygen  acids. 

Phosphonium  Compounds — (a)  Phosphonium  Iodide,  the  most  stable  of  these 
compounds,  is  obtained  by  direct  combination  of  its  components,  or,  better,  by 
adding  iodine  to  phosphorus  dissolved  in  carbon  disulphide,  distilling  off  the 
solvent,  adding  water  cautiously,  then  subliming  the  phosphonium  iodide: 

SI  +  oP  +  i6H20_»SPH4I  +  4H3P04. 

Phosphonium  iodine  occurs  in  colourless,  well-formed  crystals. 

(£)  Phosphonium  bromide  is  obtained  in  colourless  crystals  by  direct  combination 
of  phosphine  and  hydrogen  bromide  in  a  vessel  immersed  in  a  freezing-mixture. 
It  readily  dissociates  into  its  components. 

(c)  The  corresponding  chloride,  PH4C1,  is  also  obtained  in  crystalline  form  by 
direct  combination  of  its  components  in  a  freezing-mixture,  but  is  still  less  stable 
than  the  bromide,  and  dissociation  can  only  be  avoided  by  keeping  the  tempera- 
ture low,  or  by  subjecting  it  to  considerable  pressure. 

All  three  compounds  are  decomposed  by  water,  forming  the  halogen  acid  and 
phosphine,  the  latter  passing  off  as  gas. 

Liquid  Hydrogen  Phosphide,  P2H4,  is  obtained  as  a  colourless  liquid  when 
the  gas  produced  by  the  action  of  water  on  calcium  phosphide,  and  which  consists 
chiefly  of  the  hydride  PH3,  is  passed  through  a  u-tube  immersed  in  a  freezing- 
mixture.  The  liquid  boils  at  57  to  58° ;  on  exposure  to  air  it  immediately  ignites 
and  burns  with  a  brilliant  flame.  On  exposure  to  light  in  absence  of  air  it 
rapidly  decomposes  into  the  gaseous  and  solid  hydrides  : 


PHOSPHORUS  245 

Solid  Hydrogen  Phosphide,  (P4H2)3  or  Pi2H6,  obtained  as  just  described,  is 
a  yellow  flocculent  powder,  insoluble  in  water  and  alcohol.  It  is  soluble  in 
melted  yellow  phosphorus,  and  from  the  lowering  in  the  freezing-point  of  the 
solvent  (p.  196)  its  formula  has  been  proved  to  be  PiaH6.  On  heating  P12H6  in 
a  vacuum,  a  second  solid  hydrogen  phosphide,  PgHg,  is  obtained  as  a  red  powder  : 


COMPOUNDS  OF  PHOSPHORUS  WITH  THE  HALOGENS 

Phosphorus  combines  with  all  the  halogens,  forming  the  following 
compounds  — 

PF8  (gas).     PC13  (liquid).     PBr3  (liquid).     P2I4  (solid). 
PF5  (gas).     PC15  (solid).       PBr5  (solid).       PI3  (solid). 

All  of  them  can  be  obtained  by  direct  combination  of  the  elements, 
and  all  are  decomposed  by  water,  giving  rise  to  a  mixture  of  halogen 
acid  and  an  oxyacid  of  phosphorus.  This  behaviour  has  already 
been  utilized  in  the  preparation  of  hydrobromic  and  hydriodic  acids 
(pp.  156  and  161). 

Besides  the  halogen  compounds  themselves,  certain  oxyhalogen 
compounds  are  known.  The  most  important  is  phosphorus  oxy- 
chloride  or  phosphoryl  chloride,  POC13. 

Phosphorus  trifluoride,  PF3,  is  obtained  by  dropping 
arsenic  trifluoride  into  phosphorus  trichloride  : 


It  is  a  colourless  gas,  which  at  room  temperature  is  decomposed  only 
very  slowly  by  water. 

Phosphorus  pentafluoride,  PF6,  is  obtained  by  direct  com- 
bination of  the  trifluoride  and  fluorine  (Moissan)  or  by  the  action  of 
arsenic  trifluoride  on  phosphorus  pentachloride  (Thorpe)  : 


It  is  a  colourless  gas,  which  fumes  in  moist  air  and  is  decomposed 
by  water  to  form  ultimately  hydrofluoric  and  phosphoric  acids  : 


Unlike  the  corresponding  chlorine  and  bromine  compounds  (g.v.\  it 
can  be  raised  to  a  high  temperature  without  dissociating  into  the 
trifluoride  and  fluorine. 

Phosphorus  Trichloride,  PC13,  is  obtained  by  passing  dry 
chlorine  over  red  or  yellow  phosphorus  gently  heated  in  a  retort  ; 


246     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  trichloride  distils  off,  and  is  collected  in  a  cooled  receiver.  In 
order  to  remove  pentachloride",  a  little  phosphorus  is  added  and  the 
liquid  redistilled. 

Properties  —  Phosphorus  trichloride  is  a  colourless  liquid,  which 
boils  at  76°  ;  its  density  is  1.58  at  20°.  It  fumes  in  moist  air,  and  is 
at  once  decomposed  by  water  with  formation  of  phosphorous  and 
hydrochloric  acids  : 


Phosphorus  Pentachloride,  PC16,  is  best  prepared  by 
passing  dry  chlorine  into  a  large  flask  containing  the  trichloride. 
The  flask  is  kept  cool,  and  when  the  contents  become  quite  dry  it  is 
known  that  the  change  is  complete. 

Properties  —  Phosphorus  pentachloride  forms  light-yellow  almost 
colourless  crystals  with  a  pungent  odour.  When  heated  it  sublimes 
rapidly  above  140°  without  melting,  forming  a  vapour  which  is 
almost  colourless  at  low  temperatures,  but  becomes  yellower  as  the 
temperature  is  raised.  Parallel  with  this  change  in  colour  goes  a 
diminution  in  density,  both  phenomena  being  accounted  for  by 
increasing  dissociation  into  the  trichloride  and  chlorine  according  to 
the  equation 


as  already  described  (p.  169).  That  this  is  the  true  explanation  of  the 
phenomena  is  further  shown  by  the  fact  that  the  products  of  dissocia- 
tion can  be  separated  by  diffusion.  From  the  densities  it  has  been 
calculated  that  at  200°  under  atmospheric  pressure  the  pentachloride 
is  dissociated  to  the  extent  of  48.5  per  cent.,  at  250°  to  80  per  cent., 
and  at  300°  dissociation  is  practically  complete.  The  effect  of  excess 
of  chlorine  or  of  the  trichloride  in  diminishing  the  degree  of  dissocia- 
tion has  already  been  referred  to  (p.  171). 

With  a  small  quantity  of  water  the  pentachloride  is  converted  into 
phosphoryl  chloride  and  hydrochloric  acid  : 

PC15+H20->POC13+2HC1. 
With  excess  of  water  phosphoric  acid  is  formed  : 

(or  H3PO4). 


Phosphorus   pentachloride   is  largely  used,  both  in   inorganic  and 
organic  chemistry,  for  replacing  hydroxyl  groups  by  chlorine.     With 


PHOSPHORUS  247 

sulphuric  acid,  for   instance,  which  can  be  written  SO2(OH)2,  the 
following  reaction  takes  place — 

/OH  /Cl 

S02<         +  PC16->S02<(         +  POC13+HC1. 
\OH  \OH 

The  point  above  referred  to,  that  phosphorus  pentachloride  sub- 
limes before  the  melting-point  is  reached,  is  simply  a  question  of 
the  magnitude  of  the  vapour  pressure  at  the  melting-point.  It 
happens  that  before  the  melting-point  is  reached  the  vapour  pressure 
exceeds  i  atmosphere,  and  therefore  volatilization  becomes  very  rapid. 
When  heated  in  a  closed  vessel  the  pentachloride  melts  under  the 
pressure  of  its  own  vapour  at  148°. 

Phosphorus  Oxychloride,  POC13,  is  obtained,  as  mentioned 
above,  by  the  action  of  a  small  amount  of  water  on  the  pentachloride. 
It  is  a  colourless,  fuming  liquid,  which  boils  at  107°,  and  when  solid 
melts  at  1.5°.  On  treatment  with  excess  of  water  it  yields  phosphoric 
and  hydrochloric  acids,  as  stated  above. 

Phosphorus  Tribromide,  PBr3,  is  a  colourless,  fuming  liquid,  which  boils  at 
175°;  the  pentabromide  is  a  yellow  solid,  which  dissociates  into  the  tribromide 
^and  bromine  even  more  readily  than  does  the  pentachloride.  Both  bromides  are 
obtained  directly  from  their  elements,  and  they  behave  with  water  like  the  corre- 
sponding chlorides. 

Phosphorus  Diiodide,  P2I4,  is  obtained  by  adding  the  calculated  amount  of 
iodine  to  phosphorus  dissolved  in  carbon  disulphide  and  distilling  off  the  solvent. 
It  forms  orange-yellow  crystals,  which  melt  at  110°  to  a  reddish  liquid. 

Phosphorus  Triiodide,  PI3,  is  obtained  by  the  method  described  for  the  di- 
iodide  by  using  more  iodine.  It  occurs  in  dark-red  crystals,  which  melt  at  55°,  and 
are  decomposed  by  water  with  formation  of  hydriodic  and  phosphorous  acids. 


OXIDES  AND  OXYACIDS  OF  PHOSPHORUS 
Three  oxides  of  phosphorus  are  definitely  known — 

Phosphorous  oxide  (phosphorous  anhydride),     P4O6 
Phosphorus  tetroxide,  ?2O4 

Phosphorus  pentoxide  (phosphoric  anhydride),  P2O5  (P4O10) 

A  number  of  oxyacids  of  phosphorus  are  known,  most  of  which  are 
derived  from  the  oxides  just  mentioned — 

Hypophosphorousacid,H3PO2,orPH2O(OH),  )  no  corresponding 
Hypophosphoric  acid,  H2PO3  \  oxides 


248     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Phosphorous  acid,  H3PO3,  or  P(OH)3?  >  from  p  Q 

Metaphosphorous  acid,  HPO2 
Orthophosphoric  acid,  H3PO4,  or  PO(OH)3 


Pyrophosphoric  acid,  H4PoO7,  or  P2O3(OH)4 
Metaphosphoric  acid,  HPO3,  or  PO2(OH) 


from  P2O6 


Phosphorous  Oxide,  P4O6,  is  formed,  mixed  with  the  pentoxide, 
when  phosphorus  is  burned  in  a  glass  tube  in  a  current  of  dry  air.  In 
order  to  separate  the  oxides  the  products  of  combustion  are  passed 
into  the  inner  tube  of  a  Liebig's  condenser,  the  outer  tube  of  which 
contains  water  at  60°.  The  inner  tube  contains  a  plug  of  glass  wool, 
which  retains  the  pentoxide  ;  whereas  the  lower  oxide  passes  on,  and 
is  condensed  in  a  cooled  (J'tube. 

Properties  —  Phosphorous  oxide  occurs  in  colourless  crystals, 
which  melt  at  22.5°  to  a  colourless  liquid  ;  the  latter  boils  at  173.1°. 
The  oxide  is  very  poisonous.  It  combines  slowly  with  cold  water  to 
form  phosphorous  acid  : 

P406  +  6H20  =  4H3P03, 

but  when  treated  with  hot  water  a  vigorous  action  takes  place,  red 
phosphorus,  hydrogen  phosphide  and  Orthophosphoric  acid  being  the 
chief  products.  . 

When  heated  to  50-60°  in  the  air  it  inflames  and  burns  to  the  pent- 
oxide.  When  heated  in  a  sealed  tube  to  440°  it  decomposes  rapidly 
into  the  tetroxide  and  phosphorus  : 


Phosphorus  Tetroxide,  P2O4,  is  obtained  by  heating  phos- 
phorous oxide  in  a  sealed  tube  as  just  described.  It  occurs  in  colour- 
less, lustrous  crystals,  which  react  with  water  to  form  a  mixture  of 
phosphorous  and  phosphoric  acids  (cf.  N2O4,  p.  229)  : 

P204  +  3H20-»H3P03  +  H3P04. 

Phosphorus  Pentoxide,  P2O6,  is  obtained  by  burning  phos- 
phorus in  excess  of  dry  air  or  oxygen.  For  experimental  purposes  it 
is  conveniently  obtained  by  burning  phosphorus  under  a  bell  jar; 
after  a  time  the  white  fumes  settle  as  a  soft  white  powder.  It  may  be 
obtained  quite  pure  by  distilling  in  a  current  of  dry  oxygen. 

Properties  —  Phosphorus  pentoxide  is  a  colourless,  snowlike  amor- 
phous substance,  which  when  pure  is  odourless.  At  a  tempera- 
ture a  little  below  white-heat  it  is  converted  into  vapour,  which 
condenses  in  the  crystalline  form. 


PHOSPHORUS  249 

It  has  a  very  great  affinity  for  water,  and  is  the  most  efficient  drying 
agent  known  ;  its  application  for  the  complete  removal  of  moisture 
from  gases  has  already  been  referred  to.  The  first  action  of  water  is 
to  form  metaphosphoric  acid  : 

P2O5+H2O->2HPOy 

On  prolonged  standing,  much  more  rapidly  on  heating,  another  mole- 
cule of  water  is  taken  up  and  orthophosphoric  acid,  H3PO4,  is  obtained. 
An  illustration  of  the  action  of  phosphorus  pentoxide  in  abstracting 
the  elements  of  water  from  other  compounds  has  already  been  met 
with  in  the  preparation  of  nitrogen  pentoxide  : 


Vapour  density  determinations  have  shown  that  the  formula  for 
this  oxide  at  1400°  is  P4O10,  but  in  writing  equations  representing  its 
behaviour  it  is  more  convenient  to  use  the  simpler  formula. 

Hypophosphorous  Acid,  HPH2O2—  The  acid  is  most  readily 
obtained  by  the  action  of  sulphuric  acid  on  barium  hypophosphite  in 
aqueous  solution  : 

Ba(PH2O2)2+H2SO4->BaSO4  +  2HPH2O2. 

The  barium  sulphate  is  removed  by  filtration,  and  the  solution  con- 
centrated, first  in  porcelain  and  finally  in  platinum  vessels,  the  tempe- 
rature not  being  allowed  to  rise  above  105°.  On  cooling  the  acid 
separates  in  colourless  crystals. 

Barium  hypophosphite  is  obtained  in  solution  by  boiling  phos- 
phorus with  barium  hydroxide  (cf.  p.  242)  : 


Properties  —  Hypophosphorous  acid  occurs  in  colourless  leaflets* 
which  melt  at  17.4°.  On  heating  it  yields  phosphine  and  orthophos- 
phoric acid  : 

2H3PO2->H3PO4+PH3. 

Owing  to  its  tendency  to  pass  into  phosphorous  and  phosphoric  acids 
it  is  a  strong  reducing  agent.  Gold,  silver,  and  mercury  are  precipi- 
tated from  solutions  of  their  salts  : 


From  solutions  of  copper  salts,  copper  hydride,  Cu2H2,  is  precipitated 


250     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

on  warming  ;  a  reaction  which  is  characteristic  for  hypophosphorous 
acid  and  its  salts  : 


+  2Na2SO4+2Cu2H2. 

Hypophosphorous  acid  is  a  monobasic  acid,  that  is,  only  one  of  the 
hydrogen  atoms  can  be  displaced  by  metals.  The  hypophosphites 
decompose  into  phosphine  and  the  corresponding  phosphate  on 
heating.  They  are  all  soluble  in  water  ;  a  number  of  them  are  used 
in  medicine. 

The  monobasic  character  of  hypophosphorous  acid  may  be  expressed 
by  writing  its  formula  thus  : 


- 
\H 

the  assumption  being  made  that  the  acidic  properties  pertain  only  to 
the  hydrogen  attached  to  oxygen,  and  not  at  all  to  those  attached 
directly  to  phosphorus. 

Phosphorous  Acid,   H3PO3—  Preparation—  (i)  By  dissolving 
phosphorous  oxide  in  cold  water  (p.  248)  : 

P406  +  6H20->4H3P03. 

(2)  By  the  gradual  addition  of  water  to  phosphorus  trichloride,  with 
simultaneous  cooling  : 


The  product  is  then  evaporated,  the  temperature  being  finally  raised 
to  1  80°  to  drive  off  all  the  hydrogen  chloride;  on  cooling  phosphorous 
acid  crystallizes  out. 

Properties  —  Phosphorous  acid  occurs  in  colourless  crystals,  which 
'melt  at  71.1°.  On  heating  it  decomposes  into  orthophosphoric  acid 
and  phosphine  : 

4H3P03->3H3P04  +  PH3. 

It  is  a  strong  reducing  agent,  precipitating  silver  from  solutions  of 
silver  salts  and  copper  (not  copper  hydride,  difference  from  hypo- 
phosphorous  acid)  from  solutions  of  copper  salts. 

Although  it  contains  three  hydrogen  atoms,  only  two  of  them  are 
replaceable  by  metals,  and  it  is  therefore  a  dibasic  acid.  In  accord- 
ance with  this,  it  forms  two  series  of  salts.  The  two  phosphites  of 
sodium,  for  instance,  have  the  formulae  NaH2PO3  and  Na2HPO3. 


PHOSPHORUS  251 

The  behaviour  of  phosphorous  acid  may  be  expressed,  in  accordance 
with  the  principle  discussed  under  hypophosphorous  acid,  by  the 
graphic  formula 

>H 

0  =  Pf-OH 
\OH 

Metaphosphorous  Acid,  HPO2,  is  obtained  in  lustrous  crystals 
by  interaction  of  phosphine  and  oxygen  under  reduced  pressure  : 


It  dissolves  in  water  to  form  phosphorous  acid: 

+  H20->H3P03. 


Hypophosphoric  Acid,  H2PO3,  is  formed,  along  with  phos- 
phorous and  phosphoric  acids,  by  the  slow  oxidation  of  phosphorus 
in  moist  air.  The  pure  acid  occurs  in  small  colourless  crystals, 
melting  at  53°.  It  is  a  less  powerful  reducing  agent  than  phosphorous 
acid.  From  the  fact  that  it  forms  two  series  of  salts,  and  for  other 
reasons,  it  is  regarded  as  a  dibasic  acid. 

Phosphoric  Acid  (Orthophosphoric  Acid),  H3PO4—  Pre- 
paration —  (i)  On  the  large  scale,  by  the  action  of  sulphuric  acid  on 
calcium  phosphate  (bone-ash  or  natural  phosphate,  p.  239)  : 

Ca3(P04)2  +  3H2S04->3CaS04  +  2H3P04. 

The  calcium   sulphate   is   removed  by  filtration  or  by  decantation 
and  the  phosphoric  acid  concentrated  by  evaporation. 

(2)  A  purer  acid  is  obtained  by  boiling  yellow  or  red  phosphorus 
with  nitric  acid  of  density  1.2  : 

6P  +  ioHNO3  +  4H2O->6H3PO4+ioNO. 

Properties  —  Phosphoric  acid  is  usually  met  with  as  a  colourless, 
syrupy  liquid,  which,  however,  still  contains  water.  When  the  solu- 
tion is  concentrated  at  150°  and  then  allowed  to  cool,  the  pure  acid 
separates  as  colourless  crystals,  melting  at  40.7°. 

Phosphoric  acid  is  a  tribasic  acid,  forming  three  series  of  salts, 
primary  or  dihydrogen  phosphates,  e.g.  NaH2PO4,  secondary  or 
monohydrogen  phosphates,  e.g.  Na2HPQ4,  and  tertiary  or  normal 
phosphates,  e.g.  Na3PO4.  Salts  are  also  formed  from  phosphoric 
acid  by  displacing  the  hydrogen  by  different  atoms  or  groups,  thus 


252     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

sodium  ammonium  phosphate,  NaNH4HPO4,  is  well  known  as  micro- 
cosmic  salt,  and  ammonium  magnesium  phosphate,  Mg(NH4)PO4,  is 
met  with  in  analysis  (p.  464).  The  phosphates  of  the  heavy  metals 
are  mostly  normal  salts  ;  thus  the  only  known  silver  phosphate  is 
Ag3P04. 

Sodium  dihydrogen  phosphate  is  slightly  acid  in  aqueous  solu- 
tion. The  monohydrogen  phosphate,  Na2HPO4,  is,  however,  slightly 
alkaline  in  solution,  owing  to  its  being  partially  decomposed  by  water: 

Na2HP04  +  H20$NaH2P04+NaOH. 

The  alkaline  reaction  is  due  to  the  sodium  hydroxide  formed.  The 
normal  salt  is  still  more  strongly  split  up  by  water  : 

Na3PO4+HOH^:Na2HPO4 


and,  in  fact,  can  only  be  isolated  from  aqueous  solution  when  a  large 
excess  of  sodium  hydroxide  is  present.  The  hydroxide  acts,  according 
to  the  law  of  mass  action,  by  displacing  the  equilibrium  in  the  direction 
of  the  lower  arrow. 

This  decomposing  action  of  water  on  salts  is  termed  hydrolysis 
and  will  be  fully  considered  at  a  later  stage  (p.  267). 

The  normal  phosphates  of  the  alkalis  are  not  affected  by  heat,  but 
the  secondary  lose  water  and  form  pyrophosphates  : 


and  the  primary  under  similar  conditions  give  metaphosphates  : 
NaH2P04^NaP03  +  H20. 

Both  types  of  reaction  are  reversible. 

Pyrophosphoric  Acid,  H4P2O7—  Preparation—  (i)  By  heating 
orthophosphoric  acid  for  some  time  at  210  to  220°. 

(2)  In  a  purer  form  by  decomposing  the  difficultly  soluble  lead 
pyrophosphate  (prepared  by  double  decomposition  from  sodium  pyro- 
phosphate)  with  hydrogen  sulphide  : 


Pb2P207  +  2  H2S->H4 

Properties  —  The  pure  acid  occurs  in  minute  microscopic  needles 
which  melt  about  60°.  The  aqueous  solution  is  fairly  stable,  but  on 
heating  or  on  the  addition  of  traces  of  other  acids  (catalytic  action), 
it  takes  up  water  and  forms  .orthophosphoric  acid. 

Pyrophosphoric  acid  is  a  tetrabasic  acid,  but,  curiously  enough, 
forms  only  two  series  of  salts,  of  the  types  Na4P2O7  and  Na2H2P.,O7. 


PHOSPHORUS  253 

As  already  mentioned,  the  pyrophosphates  are  obtained  by  heating 
the  secondary  phosphates.  The  aqueous  solutions  are  not  affected 
even  on  boiling,  but  at  still  higher  temperatures  the  salts  take  up 
water  and  form  orthophosphates. 

Metaphosphoric  Acid,  HPO3— Preparation— (i)  By  allow- 
ing phosphorus  pentoxide  to  deliquesce  in  moist  air,  or  by  dissolving 
the  pentoxide  in  a  small  amount  of  water  : 

P206+H20-»2HP03. 

(2)  By  heating  orthophosphoric  acid  in  a  gold  crucible  at  about  400° 
till  a  molecule  of  water  is  driven  off: 

H3PO4->HPO3+H2O. 

Properties — Metaphosphoric  acid  is  a  semi-transparent  glassy  sub- 
stance, which  is  usually  sold  in  the  form  of  sticks,  and  is  known 
as  "glacial  phosphoric  acid."  It  is  readily  soluble  in  water,  and  the 
solution  is  fairly  stable  at  the  ordinary  temperature,  but  on  boiling, 
or  on  the  addition  of  a  little  mineral  acid,  it  rapidly  takes  up  water 
and  forms  orthophosphoric  acid.  This  change  can  be  detected  by 
taking  advantage  of  the  fact  that  metaphosphoric  acid  coagulates 
a  solution  of  albumen  (white  of  egg)  whereas  the  ortho  acid  does  not 
(see  below). 

Metaphosphoric  acid  is  a  fairly  strong  monobasic  acid.  The  corre- 
sponding salts,  the  metaphosphates,  can  be  prepared  directly  from 
the  acid  or  by  heating  the  dihydrogen  phosphate  (p.  252).  Many 
metaphosphates,  however,  are  complex  in  constitution,  being  most 
readily  derived  from  acids  of  the  general  formula  (HPO3)n,  where  n  is 
a  whole  number  varying  from  I  to  6  and  probably  higher.  The 
compound  Li2K4(PO3)6,4H2O,  for  instance,  is  obviously  derived  from 
the  hexabasic  acid  H6(PO3)6 ;  that  is,  n  =  6.  At  high  temperatures 
the  acid  can  be  volatilized,  and  its  vapour  density  corresponds  with 
the  formula  (HPO3)2.  At  ordinary  temperatures,  however,  it  is  doubt- 
less much  more  complex. 

Tests — The  three  phosphoric  acids  can  be  distinguished  by  their 
action  on  a  solution  of  albumen  (white  of  egg)  and  on  silver  nitrate. 
With  silver  nitrate,  orthophosphoric  acid  or  any  soluble  orthophos- 
phate  gives  a  yellow  precipitate  of  silver  orthophosphate,  Ag3PO4, 
pyrophosphoric  acid  and  pyrophosphates  a  white  precipitate  of  silver 
pyrophosphate,  Ag4P2O7,  and  metaphosphoric  acid  a  white  precipitate 
of  silver  metaphosphate,  AgPO3.  Of  the  three  acids,  only  meta- 
phosphoric coagulates  albumen.  Further,  orthophosphates  give  a 


254     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

yellow  precipitate  when  ammonium  molybdate  and  excess  of  nitric 
acid  are  added  and  the  solution  heated. 

Monoperphosphoric  Acid,  H3PO6,  and  Perphosphoric  Acid,  H4P2O8,  have  recently 
been  obtained  by  the  action  of  30  per  cent,  hydrogen  peroxide  on  phosphorus 
pentoxide  and  pyrophosphoric  acid  respectively.  They  are  powerful  oxidizing 
agents  (cf.  monopersulphuric  acid  and  persulphuric  acid  (p.  313)). 

Basicity  and  Constitution  of  Oxyacids.    Hydro  ides 

— As  already  indicated  (p.  187)  the  basicity  of  an  acid  is  measured  by 
the  number  of  hydrogen  atoms  in  a  molecule  of  the  acid  which  can  be 
displaced  by  a  metal.  The  phosphoric  acids  afford  an  excellent  illus- 
tration of  acids  of  different  basicity  obtained  by  the  combination  of  an 
acidic  oxide  with  water  in  different  proportions.  From  this  point  of 
view,  the  three  acids  may  be  represented  as  follows :  P2O5,H2O  or  2H  PO3, 
metaphosphoric  acid ;  P2O5,2H2O  or  H4P2O7,  pyrophosphoric  acid  ; 
and  P2O6,3H2O  or  2H3PO4,  orthophosphoric  acid.  Although  it  is 
sometimes  convenient  to  represent  oxyacids  as  being  formed  by  asso- 
ciation of  the  acidic  oxide  and  water  in  different  proportions,  such 
formulae  do  not  in  any  way,  and  are  not  intended  to,  represent  the 
behaviour  of  the  compound  in  question. 

The  oxyacids  may  also  be  regarded  as  hydroxides  or  derivatives  of 
hydroxides.  The  latter  are  derived  from  water  by  displacement  of 
one  of  the  H  atoms  by  another  element,  and  as  the  valencies 
of  the  different  elements  vary  from  I  to  8,  the  following  types  of 
hydroxides  are  theoretically  possible  (E  =  element). 

E(OH);  E(OH)2;  E(OH)3  ;  E(OH)4  ;  E(OH)6;  E(OH)6; 
E(OH)7;  E(OH)8. 

Hypochlorous  acid,  Cl(OH),  for  instance,  belongs  to  the  first  type, 
phosphorous  acid,  P(OH)3,  may  belong  to  the  third  type.  Further,  as 
the  same  element  may  have  different  valencies,  it  may  give  rise  to 
more  than  one  type  of  hydroxide.  The  highest  hydroxide  would  be 
that  corresponding  with  the  maximum  valency  of  the  element. 

Compounds  containing  more  than  three  OH  groups  are  usually 
unstable,  and  tend  to  form  lower  compounds  with  the  loss  of  one  or 
more  molecules  of  water.  The  normal  phosphorus  hydroxide, 
P(OH)6,  for  instance,  is  not  known,  but  the  compound  derived  from 
it  by  the  loss  of  a  molecule  of  water,  H3PO4,  is  stable. 

H0\  /OH        /OH 

\p^— OH  ->  O  —  n/  n™  '  u  ^ 
HO/  \OH 


PHOSPHORUS  255 

Similarly,  from  the  highest  nitrogen  hydroxide,  N(OH)6J  which  is 
unknown,  we  have 

N(OH)6  - 


and  from  chlorine,  the  maximum  valency  of  which  is  7,  is  derived  per- 
chloric acid,  thus — 

C1(OH)7  - 

It  has  been  suggested  that  the  hydroxyl  compound  corresponding  with 
the  maximum  valency  of  the  element  should  be  called  the  orthocom- 
pound,  thus  N(OH)5  would  be  orthonitric  acid,  and  Cl(OH)r,  ortho- 
chloric  acid.  As,  however,  many  of  these  compounds  are  unknown,  the 
convention  is  not  generally  observed.  Thus  orthophosphoric  acid, 
PO(OH)3,  is  the  first  anhydride  of  the  true  ortho  acid. 

Exactly  the  same  considerations  apply  to  bases.  The  behaviour  of 
the  hydroxide  is  determined  by  the  nature  of  the  group  E  ;  when  this 
is  a  metal,  the  hydroxide  has  basic  properties.  Compounds  are  often 
met  with  derived  not  from  the  normal  hydroxide  but  from  one  of  its 
anhydrides  ;  thus  many  bismuth  compounds  are  derived,  not  from 
Bi(OH)s,  but  from  BiO(OH),  which  is  Bi(OH)3-  H2O. 

The  connexion  between  acidic  oxides  and  acids  has  already  been 
fully  explained  (p.  186). 

The  graphic  formulas  of  the  phosphoric  acids,  which  best  represent 
their  chemical  behaviour,  are  as  follows  : — 

/OH  OH  OH    OH  OH  /OH 

0  =  PfOH;  \>  ;          0  =  P^ 

\OH  O=P-O-P=O  ^O 

Orthophosphoric  acid  Pyrophosphoric  acid  Metaphosphoric  acid 

the  phosphorus  being  quinquevaient  throughout,  and  the  basicity 
represented  by  the  number  of  hydrogen  atoms  attached  to  oxygen. 

The  graphic  formula  of  phosphorous  acid  presents  rather  more 
difficulty.  As  it  is  readily  formed  from  phosphorus  trichloride,  PC13, 

/OH 

in  which  the  phosphorus  is  trivalent,  the  formula  P^-— OH  might  be  sug- 

\OH 

gested.     As,  however,  only  two  of  the  hydrogen  atoms  can  be  displaced 

/H 
by  metals,  the  alternative  formula  O  =  P^— OH   is    preferred,   for  the 

\OH 


256     A   TEXT-BOOK    OF   INORGANIC    CHEMISTRY 

reasons  already  stated.  We  shall  meet  with  many  illustrations  of  the 
fact  that  hydrogen  atoms  attached  to  oxygen  are  more  reactive  than 
those  attached  directly  to  the  central  atom. 

The  graphic  formulas  of  the  two  remaining  oxyacids  may  be  repre- 
sented as  follows  : — 

/H 

O  =  P^— H     hypophosphorous  acid ; 
\OH 

/OH 

and  O  =  P<T          ?  hypophosphoric  acid. 
\OH 

Compounds  of  Phosphorus  and  Sulphur — We  shall  find 
at  a  later  stage  that  there  is  considerable  analogy  in  the  formulae  of 
compounds  containing  oxygen  and  sulphur  respectively  ;  due  in  part 
to  the  fact  that  both  can  function  as  bivalent  and  quadrivalent 
elements.  This  is  shown  to  some  extent  in  the  formulae  of  the 
sulphides  of  phosphorus,  which  are  obtained  by  heating  the  elements 
together  in  different  proportions,  the  following  three  being  definitely 
known,  P4S3,  P2S6  (corresponding  with  P2O6)  and  P4S7.  It  is  shown 
more  clearly  in  the  existence  of  compounds  of  the  formulas  PSC13, 
Na3PSO3  (corresponding  with  Na3PO4)  and  Na4P2S7  (corresponding 
with  Na4P2O7),  which  cannot  be  described  here. 

Phosphorus  Pentasulphide,  P2S5,  the  best-known  sulphide 
of  phosphorus,  is  obtained  by  heating  the  components,  in  the 
calculated  proportions,  in  an  atmosphere  of  carbon  dioxide.  The 
pure  compound  occurs  in  light-yellow  crystals  which  melt  at  274-276°, 
and  is  decomposed  by  water  with  formation  of  phosphoric  acid  and 
hydrogen  sulphide  : 


In  the  oxides  P2O3  (P4O6)  and  P2O5  phosphorus  is  no  doubt 
tervalent  and  quinquevalent  respectively,  though  a  little  complication 
is  introduced  by  the  fact  that  these  compounds,  from  the  results  of 
vapour  density  determinations,  have  the  formulae  P4O6  and  P4O10. 

As  P2O4  gives  a  mixture  of  phosphorus  and  phosphoric  acids  with 
water  it  presumably  contains  both  tervalent  and  quinquevalent  phos- 
phorus, its  graphic  formula  being  O=P— O— P^ 

^O 


CHAPTER  XX 

ELECTROLYSIS  AND  ELECTROLYTIC   DISSOCIATION- 
OXIDES,  ACIDS,  BASES,  AND    SALTS 

Electrical  Conductivity — Substances  which  conduct  elec- 
tricity may  be  divided  into  two  classes  :  (i)  conductors  of  the 
first  class,  in  which  the  passage  of  electricity  is  not  accompanied 
by  a  chemical  change  ;  (2)  conductors  of  the  second  class,  in  which 
the  passage  of  electricity  is  accompanied  by  chemical  decomposition. 
The  more  important  conductors  of  the  first  class  are  the  metals  and 
alloys,  but  a  few  non-metals,  such  as  graphite  (p.  310),  also  convey 
the  electric  current  without  decomposition.  For  our  present  purpose, 
however,  conductors  of  the  second  class,  the  so-called  electrolytes,  are 
more  interesting.  To  this  class  belong  aqueous  solutions  of  salts,  and 
of  so-called  "  strong "  acids  and  bases,  all  of  which  are  good  con- 
ductors. It  must  be  carefully  remembered,  however,  that  water 
itself  is  not  a  conductor,  nor  are  the  pure  salts,  acids,  or  bases 
themselves  under  ordinary  conditions.  We  have  already  seen  that 
liquefied  hydrogen  chloride  does  not  conduct  the  electric  current 
nor,  does  it  show  any  acid  properties,  but  its  solution  in  water  is 
a  good  conductor  and  also  a  strong  acid.  Although  salts  as  a  class 
under  ordinary  conditions  do  not  conduct,  it  should  be  mentioned 
that  all  of  them  readily  conduct  electricity  in  the  fused  condition, 
with  simultaneous  decomposition.  Fused  salts  form  practically  the 
only  exceptions  to  the  general  rule  that  pure  substances  (as  dis- 
tinguished from  mixtures)  are  non-conductors.  Organic  acids,  bases, 
and  salts  are  electrolytes  in  aqueous  solution,  but  practically  all 
other  organic  compounds,  whether  fused  or  in  solution,  are  non- 
electrolytes.  Finally,  it  should  be  mentioned  that  the  property  of 
electrical  conductivity  is  not  confined  to  aqueous  solutions,  but  is 
also  shown  in  a  pronounced  manner  by  solutions  in  certain  other 
solvents. 

It  has  already  been   pointed  out  that  the  decomposition  of  an 

electrolyte  when    an   electric   current  passes  is  termed  electrolysis, 

and  the  terms  employed  in  this  branch  of  the  subject  have  already 

been  mentioned.     The  plates  by  which  the  current  enters  and  leaves 

I?  357 


258     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  electrolyte  are  called  electrodes.  The  electrode  from  which 
the.  positive  current  enters  the  electrolyte  (that  connected  to  the 
positive  pole  of  the  battery  (Fig.  54))  is  called  the  anode,  the 
electrode  at  which  the  positive  current  leaves  the  electrolyte  is  the 
cathode. 

The  further  development  of  the  subject  is  best  illustrated  by  means 
of  an  example.  When  a  solution  of  sodium  sulphate,  coloured  with 
neutral  litmus  solution,  is  placed  in  a  voltameter,  and  the  electrodes 
connected  with  a  battery,  gases  are  immediately  given  off  at  the  anode 
and  cathode,  and  as  electrolysis  proceeds  the  solution  in  the  neigh- 
bourhood of  the  anode  turns  red  (showing  the  formation  of  an 
acid)  and  the  solution  round  the  cathode  turns  blue  (indicating  the 
formation  of  an  alkali).  Further  investigation  shows  that  oxygen 
only  is  given  off  at  the  anode  and  hydrogen  only  at  the  cathode. 
The  important  point  to  notice  is  that  the  chemical  changes  occur 
only  at  the  electrodes  ;  there  is  no  apparent  change  in  the  bulk  of 
solution  between  the  poles.  It  could,  however,  be  shown  by  special 
experiments  that  as  electrolysis  proceeds,  the  amount  of  sodium 
sulphate  in  the  solution  regularly  diminishes.  The  simplest  way 
of  accounting  for  these  facts  is  to  assume  that  the  dissolved  sub- 
stance is  continually  moving  towards  the  electrodes,  and  that 
chemical  changes  occur  only  when  the  materials  are  actually  in 
contact  with  the  electrodes.  This  was  the  view  taken  by  Faraday, 
who  termed  the  moving  particles  ions  (that  is,  travellers) ;  the  ions 
moving  towards  the  anode  are  called  anwns,  those  moving  towards 
the  negative  pole  or  cathode  cations. 

The  next  point  to  be  considered  is  the  nature  of  these  carriers 
in  any  particular  case.  As  regards  the  sodium  sulphate  solution, 
we  assume,  for  reasons  given  more  fully  later,  that  the  cations 
are  sodium  atoms  charged  with  positive  electricity,  while  the 
anions  are  SO4  groups  charged  with  negative  electricity.  When 
the  sodium  ions  reach  the  cathode  they  give  up  their  electrical 
charges  and  become  ordinary  metallic  sodium,  which  at  once  acts 
on  the  water,  forming  sodium  hydroxide  and  liberating  hydrogen 
according  to  the  equation 

2Na  +  2H2O  =  2NaOH  +  H2  f  (at  cathode). 

Similarly,  when  SO4  ions  reach  the  anode,  they  give  up  their 
charges  and  at  once  react  with  the  water  : 

(anode). 


ELECTROLYSIS  259 

Although  this  explanation  accounts  for  the  phenomena  just  described, 
it  is  possible  that  water  undergoes  primary  decomposition  (p.  15). 

The  electrolysis  of  sodium  sulphate  is  rather  complicated  by  the 
reaction  of  the  primary  products  of  electrolysis  with  the  solvent 
water.  A  simpler  case  would  be  the  electrolysis  of  fused  silver 
chloride,  AgCl,  with  an  anode  of  carbon  and  a  cathode  of  silver. 
In  this  case  we  may  assume  that  the  ions  are  silver  and  chlorine, 
and  during  electrolysis  silver  is  deposited  on  the  cathode  and 
chlorine  liberated  at  the  anode. 

Faraday's  Laws  —  We  will  now  consider  the  relationship 
between  the  amount  of  chemical  action  and  the  quantity  of  elec- 
tricity passed  through  a  solution.  The  amount  of  chemical  action 
might  be  estimated  by  measuring  the  volume  of  gas  liberated  at 
one  of  the  poles,  or  by  weighing  the  metal  deposited  on  an  elec- 
trode. This  question  was  investigated  by  Faraday,  who  established 
a  law  which  bears  his  name  and  which  may  be  stated  as  follows  : 
For  the  same  electrolyte^  the  amount  of  chemical  action  is  propor- 
tional to  the  quantity  of  electricity  which  passes.  Further,  Faraday 
measured  the  relative  quantities  of  substances  liberated  from  different 
electrolytes  by  the  same  quantity  of  electricity,  and  was  thus  led  to 
the  discovery  of  his  so-called  second  law  :  The  quantities  of  sub- 
stances liberated  at  the  electrodes  when  the  same  quantity  of  electricity 
is  passed  through  different  solutions  are  proportional  to  their  chemical 
equivalents.  In  a  previous  chapter  it  has  been  shown  that 

Atomic  weight=chemical        iva,ent_ 

Valency 

The  second  law  therefore  states  that  when  the  same,  quantity  of 
electricity  is  passed  through  solutions  of  such  electrolytes  as  sodium 
sulphate,  silver  nitrate,  AgNO3,  copper  sulphate,  CuSO4,  and  auric 
chloride,  AuCl3,  the  relative  amounts  of  hydrogen,  oxygen,  and 
the  metals  liberated  are  as  follows  :  — 

Electrolyte  Na2SO4  AgNO3  CuSO4      AuCl3 


The  above  law  may  also  be  stated  rather  differently  as  follows: 
The  electrochemical  equivalents  (the  proportions  of  different  elements 
set  free  by  the  same  quantity  of  electricity)  are  proportional  to  the 
chemical  equivalents. 


260     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

To  the  different  methods  of  determining  chemical  equivalents 
already  fully  discussed  (p.  126),  the  electrical  method  just  described 
has  to  be  added. 

It  is  evident  that  that  quantity  of  electricity  which  passes  through 
an  electrolyte  when  the  chemical  equivalent  of  an  element  in  grams 
is  liberated  at  each  electrode  must  be  an  important  quantity  in 
electrochemistry.  Since  I  ampere  in  i  second  (i  coulomb) 
liberates  0.00001036  grams  of  hydrogen,  it  follows  that  when  the 
chemical  equivalent  of  hydrogen  (or  any  other  element)  is  liberated 
1/0.00001036  =  96,540  coulombs  must  pass  through  the  electrolyte. 
Therefore  96,540  coulombs  will  liberate  35-45  grams  of  chlorine,  108 
grams  of  silver,  31.7  or  63.4  grams  of  copper  (according  as  the  metal 
is  in  the  cuprous  or  cupric  condition)  from  solutions  containing  the 
ions  in  question. 

Theories  of  Electrical  Conductivity — We  have  now  to 
consider  the  mechanism  of  the  conduction  of  electricity  in  an  electro- 
lyte. A  very  ingenious  theory  to  account  for  the  experimental  facts 
was  brought  forward  by  Grotthus  in  1805.  He  assumed  that  each 
molecule  is  built  up  of  positively  and  negatively  charged  particles — 
the  molecule  of  sodium  chloride,  for  instance,  of  a  positively  charged 
sodium  atom  and  a  negatively  charged  chlorine  atom — and  that  these 
molecules  are  irregularly  distributed  throughout  the  electrolyte. 
When,  however,  electrodes  in  such  a  solution  are  connected  with 
a  battery  they  exert  a  directive  effect  on  the  molecules,  the  positive 
part  of  all  the  molecules  being  turned  towards  the  negative  electrode 
and  the  negative  part  towards  the  positive  electrode.  The  sodium 
atom  next  the  cathode  then  gives  up  its  charge  to  the  latter  and 
becomes  metallic  sodium,  simultaneously  the  chlorine  half  of  another 
molecule  gives  up  its  charge  to  the  anode  and  becomes  free  chlorine. 
The  chlorine  and  sodium  atoms  thus  left  free  in  the  solution  unite 
with  the  sodium  and  chlorine  halves  of  the  neighbouring  molecules, 
so  that  an  exchange  of  partners  takes  place  all  along  the  line.  Under 
the  influence  of  the  charged  electrodes  the  molecules  then  swing 
round  and  the  process  just  described  is  repeated. 

About  1857,  however,  Clausius  pointed  out  that  the  work  done  in 
decomposing  an  electrolyte  is  entirely  expended  in  overcoming  the 
resistance  of  the  solution,  and  that  therefore  no  work  is  done  in 
pulling  apart  the  molecules  of  the  solute,  which  would  necessarily  be 
the  case  if  the  theory  of  Grotthus  were  valid.  To  get  over  this 
difficulty,  Clausius  made  the  very  important  assumption  that  under 
ordinary  conditions  a  small  proportion  of  the  molecules  of  the  solute  in 


ELECTROLYSIS  261 

an  electrolyte  are  split  up  into  their  positive  and  negative  components. 
During  their  free  moments  the  positive  and  negative  ions  would 
travel  towards  the  oppositely  charged  electrodes,  and  on  reaching 
these  would  lose  their  electrical  charges,  as  Grotthus  assumed.  The 
theory  of  Clausius  accounts  for  the  fact  that  no  work  is  done  in 
splitting  up  the  molecules,  as  the  ions  are  already  free  before 
electrolysis  is  started. 

The  Theory  of  Arrhenius.  Electrolytic  Dissociation 
—  The  view  of  Clausius  was  that  at  any  moment  only  a  minute 
fraction  of  the  molecules  were  decomposed  into  their  ions.  In  188? 
a  great  step  forward  was  made  by  Arrhenius,  who  succeeded  in  cor- 
relating the  conductance  of  solutions  with  certain  other  of  their 
properties.  We  have  seen  in  an  earlier  chapter  that  aqueous  solutions 
of  salts  and  so-called  strong  acids  and  bases  have  abnormally  high 
osmotic  pressures,  their  freezing-points  are  much  lower,  and  their 
boiling-points  much  higher,  than  those  calculated  according  to 
van  't  HofFs  theory  of  solution.  As  the  magnitude  of  the  osmotic 
pressure,  the  extent  to  which  the  freezing-point  is  depressed,  etc., 
depend  only  on  the  number  and  not  on  the  nature  of  the  particles, 
the  above  result  may  be  stated  in  the  form  that  such  solutions  behave 
as  if  they  contained  more  solute  particles  than  correspond  with  the 
ordinary  formula.  In  the  case  of  sodium  chloride,  for  instance,  the 
ratio  of  the  observed  osmotic  pressure  and  that  calculated  on  the 
assumption  that  only  NaCl  molecules  are  present,  is  about  1.7  :  i. 
Arrhenius  now  showed  that  there  is  a  close  connexion  between 
abnormally  high  osmotic  pressures  and  electrical  conductivity,  only 
those  solutions  which,  according  to  van  '/  Hoff's  theory  ',  have  abnorm- 
ally high  osmotic  pressures,  conduct  the  electric  current.  He  ascribed 
the  high  osmotic  pressures  to  a  partial  dissociation  of  the  solute  into 
charged  ions,  the  ions  acting  as  independent  particles  as  far  as  their 
effect  on  the  osmotic  pressure  and  allied  properties  is  concerned.  In 
solutions  of  sodium  chloride,  for  instance,  we  have  the  equilibrium 


and  in  normal  solution  the  salt  is  dissociated  to  about  70  per  cent. 
into  its  ions.  Similar  equilibria  occur  in  other  salts  of  the  same  type, 
for  example  — 


262     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Some  salts  may,  however,  decompose  into  more  than  two  ions.     For 
sodium  sulphate,  two  stages  of  dissociation  are  possible  : 


(i)  Na2SO4^Na  +  NaSO4.     (2) 


and  similarly  for  calcium  chloride  : 

+        — 
(i)  CaCl2^CaCl  +  Cl.    (2) 

In  the  case  of  acids  the  negative  ion  differs  in  each  case,  but  all 
give  rise  to  hydrogen  ions  : 


i  mrrr 


in  the  case  of  bases  the  positive  ion  differs  in  each  case,  but  the 
negative  ion  is  always  the  OH  group  : 

KOH^K  +  OH 

+  —  +        ++     — 

(i.)  Ca(OH)2^CaOH  +  OH.     (ii.)  CaOH^Ca  +  OH. 

It  will  be  noticed  that  in  the  above  equations  the  total  number  of 
positively  charged  ions  is  always  equal  to  the  number  of  negatively 
charged  ions.  This  is  necessarily  the  case,  as  the  solutions  are 
electrically  neutral  before  as  well  as  during  electrolysis.  As  no  atom 
has  a  smaller  charge  than  hydrogen  and  other  univalent  elements, 
we  assume  that  these  elements  have  unit  charge.  Further,  when 
sodium  sulphate  solution  is  electrolyzed  two  sodium  atoms  are  dis- 
charged for  every  SO4  group,  and  as  the  solution  remains  electrically 
neutral,  the  SO4  ion  must  have  had  two  negative  charges  to  cor- 
cespond  with  the  single  charges  of  the  two  sodium  atoms.  On  the 
same  principle  are  derived  the  other  equations  for  ionic  equilibria 
given  on  the  previous  page. 

The  application  of  the  theory  of  Arrhenius  leads  to  the  conclusion 
that  in  all  cases  the  degree  of  dissociation  increases  regularly  with 
dilution,  and  is  only  complete  when  the  dilution  is  infinitely  great. 
The  great  majority  of  salts  are  dissociated  to  more  than  50  per  cent. 
in  normal  solution,  and  those  salts  which  split  up  into  two  univalent 
ions  are  almost  completely  dissociated  in  solutions  containing  i  mol 


ELECTROLYSIS 


263 


of  salt  in  10,000  litres  of  water.  When  different  acids  are  compared, 
however,  there  is  found  to  be  a  much  greater  variation  in  the  degree 
of  dissociation.  "  Strong"  acids  are  those  which  in  ordinary  dilutions 
are  considerably  ionised  ;  that  is,  the  concentration  of  hydrogen  ions 
under  those  circumstances  is  fairly  high,  whereas  "weak"  acids,  such 
as  hypophosphorous  acid  (p.  249),  are  only  slightly  ionised  under 
ordinary  conditions,  and  their  hydrogen  ion  concentration  is  therefore 
small.  Similar  remarks  apply  to  bases.  In  aqueous  solutions  of 
"strong"  bases,  such  as  potassium  hydroxide,  the  OH'  ion1  concen- 


FIG.  54. 

tration  is  high  ;   in   solutions   of  weak  bases,  such   as  ammonium 
hydroxide,  the  OH'  ion  concentration  is  comparatively  small. 

The  mechanism  of  electrical  conductivity  on  the  Arrhenius  theory 
will  be  readily  understood  from  the  foregoing.  The  fundamental 
assumption  is  that  the  current  is  conveyed  through  the  solution  by  the 
ions  alone,  the  non-ionised  molecules  and  the  solvent  playing  no 
part  in  the  process.  When  connexion  is  made  with  the  battery  the 
ions,  which  were  previously  moving  irregularly  through  the  solution, 
are  attracted  by  the  oppositely  charged  electrodes  and  move  towards 
them  (Fig.  54).  When  they  reach  the  electrodes  they  give  up  their 
charges,  which  neutralize  a  corresponding  amount  of  electricity  on 
the  electrodes,  and  are  liberated  as  the  ordinary  uncharged  substances 

1  Instead  of  the  +  and  -  signs,  the  positive  charge  is  conveniently  indicated 
by  a  dot,  and  the  negative  charge  by  a  dash,  as  shown. 


264    A  TEXT-BOOIC   OF   INORGANIC   CHEMISTRY 

we  are  familiar  with.  In  some  cases  the  liberated  substances  appear 
as  such  (e.g.  silver  and  chlorine  when  fused  silver  chloride  is 
electrolyzed),  in  other  cases  they  attack  the  solvent  (e.g.  Na  and  SO4 
when  sodium  sulphate  solution  is  electrolyzed),  in  still  other  cases 
they  attack  the  electrodes. 

The  objection  is  sometimes  raised  to  the  ionisation  theory  that 
solutions  of  sodium  chloride,  for  instance,  cannot  contain  free  sodium 
and  free  chlorine,  because  these  would  at  once  attack  the  solvent. 
The  theory,  however,  does  not  postulate  the  existence  of  free  sodium 
and  free  chlorine  in  the  solution,  but  sodium  atoms  charged  with 
positive  electricity  and  chlorine  atoms  charged  with  negative  elec- 
tricity —  a  very  different  matter.  It  is  only  when  the  respective 
charges  are  given  up  during  electrolysis  that  the  elements  themselves 
are  liberated. 

Relationships  between  Electrical  Conductivity, 
Osmotic  Pressure,  and  other  Properties  of  Solutions 
—  Perhaps  the  most  important  feature  of  the  theory  of  Arrhenius 
is  the  establishment  of  a  quantitative  relationship  between  electrical 
conductivity  and  the  osmotic  pressure  and  allied  properties  of  electro- 
lytes. The  degree  of  dissociation  of  an  electrolyte  can  be  deduced 
quite  independently  from  electrical  conductivity  measurements,  and 
from  direct  or  indirect  measurements  of  osmotic  pressure  (most  con- 
veniently, perhaps,  by  freezing-point  measurements)  and  the  fact  that 
the  results  obtained  by  these  methods  are  in  good  agreement  is  one 
of  the  strongest  arguments  in  favour  of  the  theory. 

The  deduction  of  the  degree  of  dissociation  from  osmotic  measure- 
ments is  comparatively  simple.  Suppose  that  of  100  molecules  in  a 
solution  the  fraction  a  is  dissociated,  each  molecule  giving  rise  to  n 
ions,  the  number  of  undissociated  molecules  is  loo(i-a)  and,  the 
number  of  dissociated  molecules  being  iooa,  the  number  of  ions  is 
100  »a.  The  total  number  of  particles  is  therefore 

loo(l  -  a  +  «a)=  loo{  I  +(n  -  l)a}. 


The  ratio  of  the  number  of  particles  actually  present  to  that  calcu- 
lated according  to  Avogadro's  hypothesis  —  van't  HofPs  factor  i  —  is 
therefore 


n—  i 

We  have  seen  that  in  normal  solution  /  for  sodium  chloride  is  about 


ELECTROLYSIS  265 

1.7,  and  as  n  in  this  case  is  2,  we  obtain  a  =  0.7  ;  in  other  words, 
sodium  chloride  in  normal  solution  is  dissociated  to  the  extent  of  70 
per  cent,  into  its  ions.  As  another  illustration,  we  take  a  0.18  molar 
solution  of  calcium  chloride,  for  which  *  from  freezing-point  measure- 
ments =2.67.  As  in  this  case  the  molecule  dissociates  into  three 
ions  :  » 


«  =  3,  and,  substituting  in  the  above  formula  a=  1.67/2=0.83. 

According  to  the  theory  i  for  a  solution  of  sodium  chloride  cannot 
exceed  2,  no  matter  how  dilute  the  solution  may  be,  and,  similarly, 
the  maximum  value  of  i  for  calcium  chloride  is  3.  These  deductions 
are  fully  confirmed  by  experiment. 

It  would  lead  too  far  to  describe  fully  the  method  of  deducing  the 
value  of  a  from  conductivity  measurements,1  and  an  indication  of  the 
principle  of  the  method  must  suffice.  Other  things  being  equal,  it  is 
evident  that  the  electrical  conductivity  of  a  definite  amount  of  solute 
must,  according  to  the  theory,  be  proportional  to  the  extent  to  which 
it  is  ionised.  If  then  we  determine  the  conductivity  of  a  mol  of  salt 
at  continually  increasing  dilutions,  we  must  finally  reach  a  value  for 
the  conductivity  beyond  which  it  no  longer  increases  on  dilution. 
Under  these  circumstances  ionisation  is  complete,  in  other  words, 
the  whole  of  the  salt  is  active  in  conveying  the  current.  In  more 
concentrated  solutions  the  conductivity  under  equivalent  conditions 
must  be  less,  as  only  a  part  of  the  solute  is  active,  and  it  is  evident  that, 
other  things  being  equal,  the  ratio  of  the  conductivity  at  a  particular 
dilution  v  to  the  maximum  conductivity  is  a  measure  of  the  extent  to 
which  the  solute  is  ionised  at  the  dilution  v.  As  has  already  been 
mentioned,  the  values  of  a  deduced  by  this  method  are  in  satisfactory 
agreement  with  those  obtained  by  osmotic  methods. 

Acids  and  Bases  —  On  account  of  their  great  importance, 
acids  and  bases  require  rather  fuller  treatment  from  the  present 
point  of  view.  According  to  the  ionisation  theory  an  acid  is  a  sub- 
stance which  contains  hydrogen  ions  ;  the  characteristic  properties 
of  acids,  including  sour  taste,  action  on  litmus,  etc.,  are  due  to  these 
ions,  and  an  acid  is  the  more  active  the  greater  its  ionisation.  Thus 
normal  solutions  of  nitric  acid  and  of  nitrous  acid  contain  the  same 
concentration  of  total  hydrogen,  and  we  ascribe  the  much  greater 
activity  of  the  former  acid  to  the  fact  that  nearly  all  its  hydrogen  is 
ionised,  whereas  the  hydrogen  in  nitrous  acid  is  nearly  all  combined. 

1  Cf.  Physical  Chemistry  \  p.  261. 


266     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

The  relative  activities  of  acids  is  shown  very  clearly  by  rinding 
the  ratio  in  which  an  amount  of  base  insufficient  to  saturate  both  of 
them  distributes  itself  between  them.  If,  for  instance,  an  equivalent 
quantity  of  nitric  acid  is  added  to  a  normal  solution  of  sodium 
chloride,  an  equilibrium  is  established  : 


and  it  can  be  shown  that  about  half  the  sodium  is  present  as  the 
chloride,  the  rest  as  nitrate,  so  that  the  acids  are  approximately  of 
equal  strength.  If,  however,  hydrochloric  acid  in  equivalent  amount 
is  added  to  a  normal  solution  of  sodium  nitrite,  it  can  be  shown  that 
in  the  equilibrium 

NaNO2  +  HCl^NaCl  +  H  NO2, 


the  nitrous  acid  is  almost  completely  displaced  from  combination 
with  the  sodium,  which  is  present  almost  entirely  (97  per  cent.) 
as  sodium  chloride.  Detailed  investigation  of  such  equilibria  has 
shown  that  a  base  distributes  itself  between  two  acids  in  the  ratio 
of  their  hydrogen  ion  concentrations  as  determined  by  other 
methods  ;  a  further  justification  for  our  assumption  that  the  activity 
of  acids  as  such  is  to  be  measured,  not  by  the  total  displaceable 
hydrogen,  but  by  that  portion  of  it  which  is  present  in  the  ionic 
condition. 

It  should  not,  however,  be  as.sumed  that  all  the  properties  of  an 
acid  are  determined  by  its  H'  ion  concentration  ;  in  other  words,  that 
the  non-ionised  molecules  are  completely  inactive.  The  above  re- 
marks only  apply  to  the  properties  common  to  all  acids.  The  oxidiz- 
ing properties  of  nitric  acid,  for  instance,  are  doubtless  due,  in  part 
at  least,  to  the  non-ionised  HNO3  molecules. 

The  relative  activities  of  bases  can  also  be  compared  by  a  distribu- 
tion method  ;  that  is,  by  finding  in  what  proportions  a  quantity 
of  acid  insufficient  to  saturate  both  of  them  distributes  itself  between 
them.  In  this  way  it  may  be  shown  that  the  ratio  of  the  activities 
of  potassium  and  ammonium  hydroxides  in  normal  solution  is  about 
200  :  i,  which  is  the  ratio  of  the  OH'  ion  concentrations  in  the  two 
solutions. 

In  order  to  illustrate  the  above  remarks,  the  degree  of  dissociation, 
a,  of  a  few  acids,  bases  and  salts  at  18°  is  given  in  the  accompanying 
table.  Unless  otherwise  stated,  the  numbers  refer  to  equivalent 
normal  solutions. 


ELECTROLYSIS  267 

A  cids 

Nitric  acid        ....  0.82  Phosphoric  acid,  H'H2PO4  .     0.20 

Nitric  acid  (cone.  62  %)    .        .  0.097  Sulphurous  acid,  H'HSO3 (25°)  0.13 

Hydrochloric  acid     .         .         .  0.80  Nitrous  acid  H'NO2  (25°)      .     0.02 

Hydrochloric  acid  (cone.  62  %)  0.138  Carbonic  acid,  H'HCO3  (25°)    0.0006 

Sulphuric  acid,  H'HSO4  .        .  0.51  Hydrogen  sulphide,  H'HS  (25°)  0.00024 

Bases 

Potassium  hydroxide        .        .    0.77         Calcium  hydroxide,  N/64  at  25°  0.90 

Sodium  hydroxide    .        .         .    0.73         Silver  hy droxide,  N/4 630 at  25°  0.64 

Lithium  hydroxide  .        .        .     0.64         Ammonium  hydroxide  .         .  0.004 


Salts 

Potassium  chloride  . 
Ammonium  chloride 
Potassium  nitrate      . 
Silver  nitrate     .        . 
Sodium  acetate  . 

0.76         Potassium  sulphide 
0.75         Potassium  carbonate     . 
0.65         Barium  nitrate 
0.59         Magnesium  sulphate     . 
0.53         Copper  sulphate    . 

0.54 
0-53 
0.38 
0.26 
0.23 

.These  numbers  illustrate  the  important  fact  that  salts,  even  those 
of  weak  acids,  are  ionised  to  a  considerable  extent  in  aqueous 
solution. 

Hydrolysis  —  The  decomposing  action  of  water  on  certain  salts, 
a  process  known  as  hydrolysis,  has  already  been  referred  to,  more 
particularly  in  connexion  with  the  phosphates  of  sodium  (p.  252).  A 
much  deeper  insight  into  the  phenomena  of  hydrolysis  is  obtained  on 
the  basis  of  the  ionisation  theory. 

So  far,  water  has  been  regarded  merely  as  a  solvent,  and  the  possi- 
bility that  it  is  electrolytically  dissociated  has  not  been  taken  into 
account.  There  is,  however,  a  considerable  amount  of  evidence  to 
the  effect  that  water  itself  is  ionised,  though  to  a  very  slight  extent, 
according  to  the  equation 


so  that  it  contains  hydrogen  and  hydroxyl  ions.  The  available  evidence 
appears  to  show  that  at  room  temperature  the  concentration  of 
hydrogen  and  of  hydroxyl  ions  in  water  is  about  o.ooooooi  mol  per 
litre  ;  in  other  words,  there  is  I  mol  each  of  hydrogen  ions  and  of 
hydroxyl  ions  (i  gram  of  H'  ions  and  17  grams  of  OH'  ions)  in  ten 
million  litres  of  water.  From  the  foregoing  it  is  evident  that  water  is 
an  acid,  as  it  contains  H'  ions  ;  and  also  a  base,  as  it  contains  OH' 
ions.  These  facts  are  of  the  utmost  importance  in  connexion  with 
hydrolysis. 


t 
268     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

In  the  previous  section  it  has  been  pointed  out  that  when  two  acids 
are  allowed  to  compete  for  the  same  base  the  latter  distributes  itself 
between  the  acids  in  proportion  to  their  activities,  and  it  has  also 
been  shown  that  the  ratio  of  the  activities  of  two  acids  is  the  ratio  of 
the  extent  to  which  they  are  ionised.  The  same  considerations  apply 
for  a  salt  in  aqueous  solution  ;  water ,  in  virtue  of  its  hydrogen  ion 
concentration^  being  regarded  as  one  of  the  competing  acids.  In  the 
case  of  the  salt  of  a  strong  acid,  such  as  sodium  chloride,  it  would  not 
be  anticipated  that  such  a  weak  acid  as  water  would  take  an  appre- 
ciable amount  of  the  base,  and  the  available  experimental  evidence 
quite  bears  out  this  expectation.  In  other  words,  an  aqueous  solution 
of  sodium  chloride  contains  only  Na*  and  Cl'  ions  and  non-ionised 
sodium  chloride  in  appreciable  amount,  and  is  therefore  neutral. 

The  case  is  quite  different  for  a  salt  formed  by  a  strong  base  and  a 
weak  acid,  such  as  potassium  cyanide.  Here  water  as  an  acid  is  com- 
parable in  strength  to  hydrocyanic  acid  (p.  344),  and  there  is  therefore 
a  distribution  of  the  base  between  the  acid  and  the  water  according 
to  the  equation 

KCN  +  H2O  $  KOH  +  HCN, 

the  proportions  of  potassium  cyanide  and  potassium  hydroxide  de- 
pending upon  the  relative  strengths  of  water  and  hydrocyanic  acid 
under  the  conditions  of  the  experiment.  From  the  equation  it  is 
evident  that  potassium  hydroxide  and  hydrocyanic  acid  must  be 
present  in  equivalent  proportions,  and  since  the  hydroxide  is  much 
more  highly  ionised  than  hydrocyanic  acid  the  solution  contains  an 
excess  of  OH'  ions,  and  must  therefore  be  alkaline,  as  is  actually  the 
case. 

Analogous  considerations  apply  to  the  hydrolysis  of  salts  formed 
from  weak  bases  and  strong  acids,  such  as  ammonium  chloride. 
Although  water  is  much  weaker  than  ammonia  as  a  base,  yet  the  salt 
is  slightly  hydrolyzed  : 

NH4C1+H2O$NH4OH  +  HC1, 

and  as  hydrochloric  acid  is  much  more  highly  ionised  than  the 
hydroxide,  the  solution  contains  an  excess  of  H'  ions,  and  is  therefore 
acid.  It  should  be  remembered  that  on  hydrolysis  of  the  salt  of  a 
strong  base  and  a  weak  acid  the  solution  becomes  alkaline,  whilst  the 
solution  of  a  salt  of  a  weak  base  and  a  strong  acid  becomes  acid  on 
hydrolysis. 
As  would  be  anticipated,  salts  of  weak  bases  and  weak  acids  are 


ELECTROLYSIS  269 

hydrolyzed  to  a  still  greater  extent  under  equivalent  conditions  than 
those  belonging  to  the  types  just  considered. 

As  the'H'  and  OH7  ion  concentration  of  water  remains  practically 
constant  on  dilution,  while  the  H'  ion  concentration  of  the  com- 
peting acid  or  the  OH'  ion  concentration  of  the  competing  base 
diminishes  with  dilution,  the  hydrolysis  of  salts  of  the  first  two  types 
considered  increases  on  dilution.  Further,  as  the  ionisation  of  water 
increases  rapidly  as  the  temperature  rises,  and  the  strength  of  other 
acids  and  bases  is  only  slightly  affected  by  increase  of  temperature, 
the  hydrolysis  of  salts  is  markedly  increased  by  raising  the  tempera- 
ture. 

The  theory  of  hydrolysis  will  now  be  illustrated  by  application  to  the 
salts  of  orthophosphoric  acid,  already  considered.  It  has  been  pointed 
out  that  in  acids  containing  more  than  one  atom  of  displaceable 
hydrogen  ionisation  takes  place  in  stages,  which  in  the  present  case 
are  as  follows  :  — 

(i) 

(2) 

(3) 


Further,  in  dissociation  by  stages,  it  is  almost  invariably  the  case 
that  the  ionisation  in  the  first  stage  is  much  greater  than  in  the 
second,  and  very  much  greater  than  in  the  third  stage.  In  other 
words,  phosphoric  acid  is  a  fairly  strong  monobasic  acid,  a  much 
weaker  dibasic  acid,  and  as  a  tribasic  acid  is  very  weak  indeed.  It 
follows  that  the  salt  NaH2PO4  is  quite  stable,  the  salt  Na2HPO4  is 
slightly  hydrolyzed,  whilst  under  ordinary  conditions  the  salt  Na3PO4 
is  almost  completely  hydrolyzed  according  to  the  equation 


All  the  hydrolysis  reactions  are  reversible,  and  therefore  hydrolysis  cor- 
responds with  incomplete  neutralization,  e.g.  the  neutralization  of  potas- 
sium hydroxide  by  an  equivalent  of  hydrocyanic  acid  is  incomplete. 
Further  Applications    of   the    Ionisation   Theory  — 

The  theory  of  electrolytic  dissociation  accounts  satisfactorily  for  two 
rather  remarkable  facts  in  the  domain  of  thermochemistry.  It  has 
been  known  for  a  long  time  that  when  equivalent  amounts  of  strong 
acids  are  neutralized  by  strong  bases  in  dilute  solution  the  heat 
developed  is  the  same  in  each  case.  For  gram-equivalent  quantities  it 
amounts  to  about  13,700  calories  at  room  temperature.  Taking  as  an 
example  the  neutralization  of  sodium  hydroxide  by  hydrochloric  acid, 


270    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

and  assuming  that  ionisation  of  the  acid,  base  and  salt  is  complete, 
the  equation  may  be  written 


Na'+OH'+H-+Cl'=Na' 

Since  Na'  and  Cl'  occur  in  equivalent  amounts  on  each  side,  they 
may  be  neglected,  and  the  equation  reduces  to 

OH'+H-  =  H20, 

that  is,  the  combination  of  H*  and  OH'  ions  to  form  water.  The 
same  equation  applies  to  the  neutralization  of  any  other  strong  base 
by  a  strong  acid  ;  provided  that  the  solutions  are  so  dilute  that  ionisa- 
tion is  practically  complete,  the  process  in  all  cases  consists  simply  of 
the  combination  of  H*  and  OH'  ions  to  form  non-ionised  water.  We 
may  therefore  anticipate  that  the  heat  development  will  be  the  same 
for  equivalent  quantities,  as  is  actually  the  case. 

The  next  point  to  consider  is  the  so-called  law  of  thermo-neutrality^ 
that  no  heat  change  occurs  on  mixing  dilute  solutions  of  two  highly 
dissociated  salts  ;  for  example,  sodium  chloride  and  potassium  nitrate. 
According  to  the  ionisation  theory,  assuming  that  dissociation  is  prac- 
tically complete,  the  reaction  may  be  represented  as  follows  :  — 


The  right  and  left  hand  sides  of  the  equation  are  identical,  so  that 
the  salts  exert  no  mutual  influence,  which  accounts  for  the  fact  that 
no  thermal  change  occurs  on  mixing. 

The  law  of  mass  action  applies  to  electrolytes  (with  some  ex- 
ceptions) just  as  to  equilibria  in  which  undissociated  chemical  com- 
pounds are  concerned.  For  instance,  the  addition  of  excess  of  one  of 
the  products  of  ionisation  to  an  electrolyte  diminishes  the  ionisation, 
just  as  excess  of  one  of  the  products  drives  back  ordinary  dissociation 
(p.  171).  Thus  the  strength  of  a  weak  acid,  such  as  nitrous  acid,  is 
greatly  diminished  by  adding  excess  of  a  highly  dissociated  nitrite  : 

HN02;±H-  +  N02', 

the  equilibrium  being  driven  towards  the  left.  Also  the  strength  of 
ammonia  is  greatly  diminished  by  addition  of  an  ammonium  salt  such 
as  the  chloride,  and  advantage  is  often  taken  of  these  facts  in  ana- 
lytical chemistry. 

In  the  foregoing  we  have  confined  our  attention  almost  entirely  to 
electrolytic  dissociation  in  aqueous  solutions,  and  this  is  justified  by 


ELECTROLYSIS  271 

the  predominant  use  of  water  as  a  solvent  in  chemical  operations.  It 
should  be  mentioned,  however,  that  other  solvents  have  also  been 
extensively  investigated  from  this  point  of  view,  and  it  has  been  found 
that  many  of  them  are  practically  incapable  of  forming  conducting 
solutions;  whilst  others,  such  as  liquefied  ammonia  (p.  215),  form 
highly  conducting  solutions  with  many  solutes. 

Electrolytic  Dissociation  and  Valency — The  electro- 
lytic dissociation  theory  throws  some  light  on  the  valency  of 
electrolytes.  According  to  the  views  just  developed,  the  metallic 
components  of  salts  in  solution  are  looked  upon  as  being  positively 
charged,  the  number  of  charges  corresponding  with  the  ordinary 
valencies  of  the  metals.  Some  important  cations  are  K*,  Na",  Ag', 
NH4*,  Ca",  Kg",  Fe",  Al'",  etc.  The  remainder  of  the  salt  mole- 
cule constitutes  the  negative  ion  (anion)  which,  like  the  positive  ion, 
may  have  one,  two  or  more  (negative)  electric  charges  corresponding 
with  its  ordinary  valency.  Among  the  more  important  anions,  the 
majority  of  which  have  already  been  mentioned,  are  Cl',  Br',  I',  NO3', 
CICy,  SO/,  CO3",  PO4'",  etc.  The  nature  of  the  ions  is  of  course 
determined  by  examining  the  products  of  electrolysis  at  the  anode 
and  cathode  respectively.  It  appears  from  the  above  that  certain 
elements,  more  particularly  hydrogen  and  the  metals,  have  an  affinity 
for  positive  electricity,  and  never  appear  alone  in  solution  otherwise 
than  as  positive  ions,  whilst  certain  other  atoms  and  groups  have 
an  affinity  for  negative  electricity  and  occur  in  solutions  only  as 
negative  ions. 

It  should,  however,  be  remembered  that  though  metals  never 
occur  alone  in  the  ionised  condition  otherwise  than  as  cations, 
they  may  form  part  of  complex  anions,  and,  similarly,  elements 
which  alone  appear  only  as  negative  ions  may  form  part  of  com- 
plex anions.  For  example,  the  first  stage  in  the  dissociation  of 
calcium  chloride  is 

CaCL^CaCl'+Cl', 

one  of  the  chlorine  atoms  being  present  in  the  positive  ion.  It 
should  be  emphasized  that  the  ordinary  tests  for  the  elements 
apply  to  them  only  when  occurring  as  simple  ions,  thus  KC1O3  or 
KC1O4  give  none  of  the  tests  for  chlorides,  which  are  given  only 
by  the  Cl'  ion  itself. 

With  regard  to  the  non-ionised  salts  it  is  natural  to  assume, 
although  there  is  no  direct  evidence  on  the  point,  that  the  atoms 
are  held  together,  at  least  partly,  by  the  attraction  of  the  contrary 


272     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

electrical  charges.      The   sodium   chloride   molecule,   for    instance, 

+  - 
may  be  formulated  thus,  NaCl. 

It  is  clear,  then,  that  salts  (including  acids  and  bases)  contain 
atoms  or  groups  with  contrary  electrical  charges,  but  there  is  no 
evidence  that  this  is  true  of  the  enormous  number  of  chemical 
compounds  which  do  not  undergo  electrolytic  dissociation.  We 
have  seen,  for  instance,  that  chlorine  and  iodine,  which  are  re- 
garded as  electro-negative  elements,  unite  atom  for  atom  to  form 
the  well-defined  compound  IC1.  Of  the  nature  of  the  chemical 
attraction  in  this  case,  whether  electrical  or  otherwise,  we  know 
nothing.  The  same  remark  applies  to  the  molecules  of  the  elements 
themselves.  In  some  instances  (e.g.  hydrogen,  oxygen)  these  are 
so  stable  that  even  at  the  highest  temperatures  at  our  disposal 
there  is  no  evidence  even  of  a  partial  decomposition  into  the  com- 
ponent atoms. 

Finally,  while  on  the  subject  of  valency,  a  few  words  may  be 
devoted  to  compounds  formed  by  the  association  of  two  or  more 
complete  molecules,  e.g.  HF'KF,  and  NaCl,2H2O.  Substances  of 
this  type,  of  which  the  commonest  are  salts  associated  with  "  water 
of  crystallization,"  are  often  called  "  molecular  compounds,"  and 
cannot  be  represented  satisfactorily  on  the  simple  theory  of  valency. 
If  the  affinities  of  the  atoms  of  sodium  and  chlorine  were  completely 
satisfied  in  forming  sodium  chloride,  it  is  scarcely  to  be  anticipated 
that  the  molecule  could  bind,  though  weakly,  two  molecules  of  water. 
We  may  assume,  then,  that  there  remains  over  a  little  affinity,  the 
so-called  "  free  affinity,"  which  enables  the  molecule  to  bind  one  or 
more  molecules  of  another  compound  which  also  possesses  free 
affinity.  Few,  if  any,  inorganic  compounds  appear  to  be  entirely 
devoid  of  residual  affinity  or  entirely  saturated ;  in  other  words,  all 
of  them  appear  to  be  capable  of  forming  more  or  less  stable  "mole- 
cular compounds  "  with  other  compounds.  Some  substances  belong- 
ing to  organic  chemistry,  however  (p.  347),  appear  to  be  almost 
completely  saturated. 

At  present,  the  subject  of  valency  is  in  a  rather  unsatisfactory 
condition,  and  the  views  of  chemists  differ  widely  on  many  points 
of  the  subject.  The  matter  is  further  referred  to  on  p.  578. 


OXIDES,    ACIDS,    BASES,   AND   SALTS  273 

OXIDES,   ACIDS,   BASES,   AND    SALTS 

We  are  now  in  a  position  to  consider  more  fully  the  nature  of  the 
more  important  classes  of  chemical  compounds,  such  as  oxides,  acids, 
bases,  and  salts,  and  to  trace  briefly  the  stages  by  means  of  which 
our  present  conceptions  as  to  these  classes  of  compounds  have  been 
reached,  Our  discussion  will  serve  at  the  same  time  as  a  summary 
of  what  has  already  been  stated  regarding  such  compounds. 

Classification  of  Oxides— After  the  true  nature  of  combustion 
had  been  established  by  Lavoisier  much  attention  was  devoted  to 
the  compounds  of  oxygen  with  other  elements,  and  it  was  soon 
recognized  that  the  oxides  differ  greatly  among  themselves  in 
chemical  behaviour  and  must  be  placed  in  different  classes.  Owing 
to  this  diversity  in  chemical  behaviour  it  is  difficult  even  at  the 
present  day  to  give  a  satisfactory  classification  of  oxides.  The 
following  main  groups  may  be  distinguished  : — 

(1)  Acidic  Oxides  or  Anhydrides,  which  combine  with  water  to 
form  acids  and  react  with  bases  to  form  salts.     Example  :  Chlorine 
monoxide,  C12O.     The  majority  of  acidic  oxides  are  oxygen   com- 
pounds of  non-metals. 

(2)  Basic  Oxides,  which  combine  with  water  to  form  bases  and 
react  with  acids  to  form  salts.     Example  :  Calcium  oxide,  CaO.    The 
majority  of  basic  oxides  are  oxygen  compounds  of  metals. 

(3)  True  Peroxides,  which  contain  a  higher  proportion  of  oxygen 
than  the  basic   oxide   of  the  metal   (corresponding  with   its   usual 
valency)   and  yield  hydrogen    peroxide    on    treatment    with    acids. 
Example  :  Barium  peroxide,  BaO2  : 

BaO2  -f  H2SO4->BaSO4  +  H2O2. 

(4)  Polybxides,  which  contain  a  higher  proportion  of  oxygen  than 
the  basic  oxide  of  the  metal,  and  with  concentrated  acids  yield  salts 
corresponding  with  the  normal  oxide,  and,  in  addition,  either  oxygen 
or  a  product  of  oxidation  of  the  acid.     Example  :  MnO2  (p.  88)  : 

MnO2 +4HCl->MnCl2  +  C12  +  2H2O. 

(5)  Neutral  or  indifferent  Oxides,  which  do  not  belong  to  any  of 
the  above  classes.     Example  :  Nitric  oxide,  NO  ;  water,  H2O. 

There  are  certain  oxides  which  may  be  classed  both  as  basic  and 
acidic,  since  they  give  salts  both  with  acids  and  bases,  e.g.  aluminium 
oxide,  A12O3.     They  are  sometimes  termed  amphoteric  oxides  (ampho- 
18 


274     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

teros  =  both).  Lead  dioxide,  PbO2,  is  not  only  an  amphoteric  oxide 
but  often  behaves  as  a  polyoxide,  and  therefore  it  belongs  to  no  less 
than  three  of  the  above  classes  (p.  504). 

Peroxides  are,  in  a  sense,  polyoxides,  but  differ  from  the  other 
members  of  this  class  in  giving  hydrogen  peroxide  with  acids.  It  is 
generally  assumed  that  the  valency  of  the  metal  in  the  peroxides  is 

/O 
the  same  as  that  in  the  normal  oxides,  e.%.  Ba  =  O  ;  Ba^    |  ,  whereas 

\0 
in  the  polyoxides  the  metal  functions  with  a  higher  valency  than  in 

//® 
the  basic  oxide,  e.g.  Pb  =  O  ;  Pbx'     .     There  is  independent  evidence 


that  lead  can  behave  as  a  quadrivalent  element  (cf.  p.  507). 

Some  oxides  appear  to  be  formed  by  combination  of  basic  and 
acidic  oxides  of  the  same  metal,  and  are  of  the  nature  of  salts, 
e.g.  Pb3O4  may  be  regarded  as  2PbO,PbO2. 

Preparation  of  Oxides—  A  number  of  different  methods  of 
preparing  oxides  have  already  been  given  in  the  foregoing  pages. 
The  more  important  general  methods  are  as  follows  : 

(1)  Direct  combination  of  the  element  ivith  oxygen.     Examples: 
Lead  oxide,  PbO  ;  Nitric  oxide,  NO  ;  Phosphorus  pentoxide,  P2O6. 

(2)  Heating  another  oxide   in   air  or  oxygen  —  e.g.    trimanganic 
tetroxide,  Mn3O4  ;  red  lead,  Pb3O4. 

(3)  Heating  the  oxyacid  (for  acidic  oxides}  or  the  hydroxide,  car- 
bonate^  or  nitrate  of  a  metal  (for  basic  oxides}.     Examples  : 

2HI03=I206+H20, 
.BaCO3=BaO-»-CO2. 

(4)  Action  of  an  oxidizing  agent  (nitric  acid,  chlorine,  &c.)  on  the 
element  or  on  a  lower  oxide  : 

Sn  +  2HNO->SnO 


Besides  the  above  general  methods  a  number  of  special  methods 
are  also  used. 

The  Nature  of  Acids—  We  have  already  seen  (p.  26)  that  the 
oxides  of  non-metals  are  usually  acidic,  combining  with  water  to  form 
acids.  This  observation  led  Lavoisier  to  the  view  that  acids  are  the 
oxides  of  non-metallic  elements,  oxygen  being  the  acidifying  principle 
common  to  all  acids.  As  metals  on  combination  with  oxygen  gave 


OXIDES,    ACIDS,    BASES,   AND    SALTS  275 

basic  oxides  which  formed  salts  with  acids,  it  was  assumed  that  salts 
were  formed  by  combination  of  acidic  and  basic  oxides.  On  this  view 
calcium  sulphate  is  represented  as  CaO,SO3  and  normal  calcium 
phosphate  as  3CaO,P2O6  (cf.  p.  453)- 

The  Oxygen  Theory  of  Acids^  as  it  was  called,  was  accepted  by 
nearly  all  chemists  for  a  good  many  years  after  the  death  of  Lavoisier, 
although  certain  facts  were  known  in  apparent  contradiction  with  it. 
Thus  Berthollet  in  1787  found  that  hydrocyanic  acid  contained  only 
hydrogen,  carbon,  and  nitrogen  (p.  344) ;  and  it  was  known  to  him 
that  hydrogen  sulphide,  which  has  distinct  acidic  properties,  con- 
tained only  hydrogen  and  sulphur  (p.  293).  The  question  which  led 
to  the  final  overthrow  of  the  oxygen  theory  of  acids  was  the  constitu- 
tion of  chlorine  and  hydrochloric  acid.  In  accordance  with  the 
oxygen  theory  of  acids,  hydrochloric  acid  (muriatic  acid)  was  regarded 
by  Gay-Lussac  and  Thenard  as  a  compound  of  oxygen  with  a  hypo- 
thetical substance,  muriaticum  (and  water),  whilst  chlorine,  which  is 
formed  from  hydrochloric  acid  by  oxidation,  was  supposed  to  contain 
either  more  oxygen  or  less  hydrogen  than  the  latter.  Sir  Humphry 
Davy  (1810)  showed,  however,  that  no  oxide  of  carbon  is  formed 
when  carbon  is  raised  to  a  white  heat  in  oxymuriatic  acid  (chlorine), 
that  the  gas  is  not  decomposed  by  the  prolonged  action  of  a  powerful 
electric  discharge,  and  that  the  compounds  formed  by  the  action  of 
oxymuriatic  acid  on  phosphorus  and  tin  contained  no  oxygen.  Gay- 
Lussac  and  Thenard  themselves  were  unable  to  obtain  any  oxide  of 
carbon  by  heating  carbon  in  hydrogen  chloride.  Davy's  experiments 
proved  conclusively  that  chlorine  is  an  element  and  that  hydrogen 
chloride  contains  no  oxygen,  so  that  the  oxygen  theory  of  acids 
became  untenable. 

The  Hydrogen  Theory  of  Acids — The  Swedish  chemist  Berzelius, 
who  was  the  chief  authority  on  chemical  questions  at  this  period,  for  a 
long  time  opposed  the  new  views  as  to  the  constitution  of  hydro- 
chloric acid  because  they  were  opposed  to  his  dualistic  theory  of 
chemical  compounds.  According  to  this  theory,  which  is  considered 
more  fully  under  salts  (p.  277),  salts  are  regarded  as  compounds  of 
acidic  and  basic  oxides  and  acids  as  compounds  of  acidic  oxides 
with  water,  the  latter  acting  as  a  weakly  basic  constituent.  On  this 
view  the  formula  of  sulphuric  acid  is  H2O,SO3,  and  of  nitric  acid 
N2O5,H2O.  When  the  evidence  that  certain  acids  contain  no  oxygen 
could  no  longer  be  gainsaid,  Berzelius  drew  a  sharp  distinction 
between  haloid  acids,  compounds  of  halogen  and  hydrogen,  and 
oxyacids,  which  contain  oxygen. 


276     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

As  early  as  1815  Davy,  who  had  observed  that  iodic  anhydride, 
I2O6,  has  itself  no  acidic  properties  but  acquires  them  only  after 
combination  with  water,  expressed  the  view  that  hydrogen,  and  not 
oxygen,  is  the  essential  constituent  of  all  acids.  Somewhat  similar 
views  were  expressed  by  Dulong.  At  the  time,  however,  the  electro- 
chemical theory  of  Berzelius  held  the  field  and  Davy's  views  found  no 
acceptance. 

To  this  period  belong  the  important  investigations  of  Graham  on 
the  phosphoric  acids  which  led  to  the  enunciation 'of  the  theory  of 
polybasic  acids.  At  that  time  it  was  assumed  that  the  "atoms" 
(molecules)  of  most  metallic  oxides  contained  one  atom  of  the  metal 
to  one  atom  of  oxygen,  and  combined  with  one  "  atom  "  (molecule)  of 
acid  to  form  neutral  salts.  On  this  view  practically  every  acid  was 
regarded  as  monobasic.  Graham  now  showed  that  ortho-,  pyro-,  and 
metaphosphoric  acids  contained  3,  2,  and  I  molecules  of  "  basic  water" 
in  combination  with  one  molecule  of  P2O6.  These  molecules  of  water 
could  be  displaced  by  metallic  ,  oxides  with  formation  of  salts.  It 
followed  that  the  combining  capacity  of  the  phosphoric  acids  for 
bases  depended  on  the  amount  of  water  associated  with  the  acidic 
oxide.  As  a  result  of  these  and  of  his  own  investigations  Liebig  was 
led  to  enunciate  the  theory  of  polybasic  acids,  according  to  which 
acids  differ  in  their  combining  capacity  for  bases.  Liebig  used  as  a 
criterion  for  polybasic  acids  their  capacity  to  form  compound  salts 
with  different  metallic  oxides;  for  example,  the  salt  Na2KPO4  (or 
2Na2O,K20,P2O6). 

Both  Graham  and  Liebig  interpreted  their  results  on  the  basis  of 
the  electrochemical  theory  of  Berzelius,  but  at  a  later  stage  Liebig 
was  led  to  reintroduce  the  theory  of  hydrogen  acids  originally  brought 
forward  by  Davy  and  Dulong.  One  of  the  chief  reasons  which  led 
him  to  this  opinion  was  that  the  oxygen  acids  and  haloid  acids  could 
then  be  regarded  from  the  same  standpoint,  a  conclusion  to  which  one 
is  practically  compelled  by  their  very  similar  behaviour.  Liebig  stated 
his  views,  which  are  valid  at  the  present  day,  as  follows  : 

(1)  Acids  are  particular  compounds  of  hydrogen,  in  which  the  latter 
element  can  be  displaced  by  metals. 

(2)  Neutral  salts  are  those  compounds  of  the  same  class,  in  which 
the  hydrogen  has  been  displaced  by  its  equivalent  of  a  metal.     Those 
substances  which  we  at  present  term  anhydrous  acids  for  the  most 
part  only  acquire  their  property  of  forming  salts  with  metallic  oxides 
after  the  addition  of  water,  or  they  are  compounds  which  decompose 
these  oxides  at  fairly  high  temperatures. 


OXIDES,   ACIDS,    BASES,    AND   SALTS  277 

The  last  sentence  is  worthy  of  note.  We  have  already  seen  that 
liquefied  hydrogen  chloride  has  no  acid  properties  (p.  96) ;  it  only 
acquires  such  on  addition  of  water. 

Since  Liebig's  time  new  light  has  been  thrown  on  the  nature  of 
acids  on  the  basis  of  the  electrolytic  dissociation  theory,  as  has  been 
fully  explained  in  the  course  of  the  chapter  (p.  261).  According  to 
this  theory,  acidic  properties  are  due  to  the  presence  of  hydrogen 
ions,  and  the  strength  of  an  acid  is  determined  by  the  extent  to  which 
it  is  ionised  under  the  conditions  of  experiment.  The  inactivity  of 
liquefied  hydrogen  chloride  is  ascribed  to  the  absence  of  hydrogen 
ions  ;  it  becomes  ionised  only  on  addition  of  water. 

The  Nature  of  Bases— The  development  of  our  views  on 
bases  has  been  referred  to  incidentally  in  connection  with  acids,  and 
can  therefore  be  dismissed  quite  briefly.  The  oxides  of  metals,  which 
with  acids  form  salts,  are  termed  basic  oxides,  and  on  addition  of 
water  they  form  bases.  According  to  the  dualistic  theory,  the  water 
in  bases  was  assumed  to  play  the  part  of  a  weak  acidic  oxide,  and 
potassium  hydroxide,  for  instance,  could  be  formulated  thus,  K2O, 
H2O.  When  the  dualistic  theory  was  overthrown,  the  simpler 
unitary  formulas,  such  as  KOH  and  Ca(OH)2,  took  the  place  of  the 
previous  formulas.  According  to  the  ionisation  theory,  also  in  a 
sense  a  dualistic  theory  (p.  261),  bases  are  ionised  into  the  negative 
OH'  group  and  a  positive  group,  and  the  presence  of  the  OH'  group 
is  just  as  characteristic  of  bases  as  the  presence  of  H  ions  is  of  acids. 

Just  as  there  are  polybasic  acids,  so  polyacidic  bases  are  known, 
e.g.  Ca(OH)2 ;  Bi(OH)3,  from  which  salts  can  be  derived  by  the 
partial  or  complete  displacement  of  the  hydroxyl  groups  by  electro- 
negative groups  (p.  281). 

The  Nature  of  Salts— As  already  mentioned,  Lavoisier,  the 
founder  of  modern  chemistry,  regarded  salts  as  being  formed  by  the 
combination  of  an  acid  and  a  base,  or,  as  we  would  say  at  the  present 
day,  by  combination  of  an  acidic  and  a  basic  oxide.  This  view  was 
later  developed  by  Berzelius,  mainly  on  the  basis  of  electro-chemical 
evidence,  into  the  so-called  dualistic  theory,  which  for  many  years 
exerted  a  great  influence  on  the  development  of  chemistry. 

In  1803  it  was  shown  by  Berzelius  that  when  salts,  acids,  &c.,  are 
subjected  to  the  action  of  an  electric  current,  hydrogen,  the  metals, 
metallic  oxides,  &c.,  are  liberated  at  the  negative  pole,  whilst  oxygen, 
acids,  &c.,  are  given  off  at  the  positive  pole.  Thus,  when  a  solution 
of  sodium  sulphate  is  electrolysed,  sodium  hydroxide  is  obtained  at 
the  negative  pole  and  sulphuric  acid  at  the  positive  pole  (p.  258). 


278     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

From  these  and  other  considerations  Berzelius  was  led  to  the  view 
that  all  substances  consist  of  an  electro-positive  and  an  electro- 
negative part,  held  together  by  electrical  attraction.  Thus  salts  are 
formed  by  the  combination  of  basic  oxides,  which  are  electro-positive, 

+ 

and  acidic  oxides,  which  are  electro-negative,  e.g.  Na2O,  SO3,  acids 
by  the  combination  of  acidic  oxides  with  water  acting  as  a  weak  base> 

+ 
e.g.  H2O,  SO3,  and  bases  by  a  combination  of  basic  oxides  with  water 

+ 

acting  as  an  oxide  with  weak  acidic  properties,  e.g.  CaO,  H2O. 
These  groups  themselves  were  regarded  as  being  formed  of  elements 

+  - 
of  opposite  polarity,  thus  CaO  has  an  excess  of  positive  electricity 

I      

and  SO3  an  excess  of  negative  electricity.  When  such  groups  unite 
to  form  salts  a  partial  neutralization  of  the  electrical  charges  occurs. 
The  atoms  themselves  were  regarded  as  bipolar  in  character,  one  or 
other  of  the  charges  being  in  excess. 

The  fact  established  by  Davy  that  chlorine  is  an  element,  and  not 
"oxygenated  muriatic  acid,"  to  some  extent  shook  confidence  in  the 
dualistic  theory,  although  the  majority  of  chemists  continued  to  hold 
it.  The  reasons  which  led  to  the  abandonment  of  the  binary  theory 
belong  mainly  to  the  domain  of  organic  chemistry  and  can  only 
be  touched  upon  here.  For  instance,  hydrogen  was  regarded  by 
Berzelius  as  a  positive  element  and  chlorine  as  a  negative  element, 
and  it  was  to  be  anticipated  that  the  substitution  of  chlorine  for 
hydrogen  in  a  compound  would  entirely  alter  its  character.  As  a 
matter  of  fact,  it  was  shown  by  Dumas  that  the  compound  formed 
by  the  displacement  of  three  hydrogen  atoms  in  acetic  acid  by 
chlorine  is  still  a  monqbasic  acid,  and  in  most  respects  behaves 
very  like  acetic  acid  itself.  Further,  the  compounds  potassium 
permanganate,  KMnO4,  and  potassium  perchlorate,  KC1O4,  which 
differ  only  in  that  in  the  latter  an  atom  of  chlorine  has  taken  the 
place  of  an  atom  of  manganese  in  the  former,  are  both  salts  of 
monobasic  acids  and  are  isomorphous,  although  no  two  elements 
differ  more  than  chlorine  and  manganese.  Facts  like  these  led 
ultimately  to  the  abandonment  of  the  dualistic  theory  and  to  the 
acceptance  of  a  unitary  theory  of  chemical  compounds,  according 
to  which  a  chemical  compound  must  be  regarded  as  a  whole ;  its 
properties  depending  primarily  upon  the  arrangement  and  number 


OXIDES,    ACIDS,    BASES,   AND   SALTS  279 

of  the  atoms,  and  in  a  lesser  degree  on  their  nature.  Later  investiga- 
tion has  shown  that  Dumas  undervalued  the  effect  of  the  nature  of 
the  constituent  atoms  on  the  properties  of  chemical  compounds. 

About  this  time  the  radical  theory  and,  at  a  later  stage,  the  theory 
of  types  played  a  great  part  in  the  development  of  chemistry,  mainly 
organic  chemistry.  According  to  the  latter  theory,  the  majority  of 
chemical  compounds  then  known  could  be  regarded  as  belonging  to 
one  of  four  types,  the  hydrogen,  water,  ammonia,  and  methane  types, 
formulated  as  follows  : 


HI    Hi      HI     H! 

\0        HJN  [ 

H)         HJ  HJ  H 


Chemical  compounds  were  looked  upon  as  being  derived  from  one 
of  these  four  types  by  the  complete  or  partial  substitution  of  their 
hydrogen  atoms  by  residues  or  radicals.  Thus  hydrogen  chloride 

and  ethyl  chloride  belong  to  the  hydrogen  type  ri  f  i  2  Cl  1" 
potassium  hydroxide  and  nitric  acid  belong  to  the  water  type,  being 
formulated  thus  :  TT  [  O  ;  -J  [•  and  so  on.  A't  a  later  stage  it  was 

found  necessary  to  introduce  mixed  types  and  the  matter  became 
rather  complicated. 

Although  the  so-called  type  theory  is  more  a  system  of  classifi- 
cation than  a  theory  of  chemical  compounds,  there  is  a  principle  of 
the  greatest  importance  underlying  it,  which  we  now  proceed  to 
consider.  In  1852  Frankland  pointed  out  that  nitrogen,  phosphorus, 
arsenic,  and  antimony  tend  to  form  compounds  containing  for  each 
atom  of  one  of  these  elements  either  three  or  five  equivalents  of  other 
elements.  Of  compounds  constituted  according  to  the  ratio  i  to  3 
we  have  NO3,  NH3,  NI3,  PO3,  PC13,  SbO3,  AsH3,  AsO3,  &c.,  and 
according  to  the  ratio  i  to  5  the  compounds  NO6,  NH4I,  PO6,  &c. 
Frankland  proceeds  as  follows :  Without  putting  forward  any  hypo- 
thesis as  to  the  reason  of  this  agreement  in  the  grouping  of  the  atoms 
it  is  evident  that  such  a  tendency  or  regularity  exists,  and  that  the 
affinity  of  an  atom  of  one  of  these  four  elements  is  always  satisfied  by 
the  same  number  of  atoms  (equivalents}  of  other  elements  ivithout 
regard  to  the  chemical  character  of  the  latter.  This  is  the  first  state- 
ment of  the  views  from  which  the  theory  of  valency  was  developed  by 
Frankland  himself,  Kekule,  Couper,  and  others  (cf.  p.  130).  It  will 
now  be  clear  that  the  form  of  the  different  types  is  determined  by  the 


280     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

combining  capacity  or  valency  of  the  elements  present.  Further, 
potassium  hydroxide  and  nitric  acid  belong  to  the  water  type  because 
the  K  and  NO2  groups  are  univalent  and  therefore  can  take  the  place 
of  one  atom  of  hydrogen.  The  further  development  of  the  theory  of 
valency  led  to  the  representation  of  chemical  compounds  by  means 
of  graphic  formulas,  which  show  the  mode  of  linking  of  the  atoms  in 
the  molecule,  and  are  meant  to  represent  as  far  as  possible  the 
behaviour  of  chemical  compounds.  We  have,  however,  already  met 
with  a  great  deal  of  evidence  showing  that  graphic  formulae  are  at 
best  a  very  unsatisfactory  representation  of  the  many-sided  behaviour 
of  the  majority  of  chemical  compounds. 

Salts  and  the  lonisation  Theory— Since  1887  a  new  theory 
as  to  the  nature  of  solution  of  salts,  acids,  and  bases — the  electrolytic 
dissociation  theory— has  secured  almost  universal  acceptance.  This 
theory,  and  the  evidence  on  which  it  is  based,  have  been  fully 
described  in  the  course  of  this  chapter.  In  this  connection  two  points 
are  worthy  of  mention.  In  the  first  place  the  theory,  like  the  dualistic 
theory  of  Berzelius,  is  based  largely  on  electro-chemical  evidence,  and 
further,  it  is  dualistic  in  character.  There  is,  however,  a  fundamental 
difference  between  ^he  two  theories,  inasmuch  as  according  to  the 
ionisation  theory  the  two  constituents  are  the  metal  or  a  group 
behaving  like  a  metal  (positively  charged)  and  the  remainder  of 
the  molecule  (negatively  charged),  whilst  according  to  the  view  of 
Berzelius  a  salt  is  formed  by  the  combination  of  a  basic  with  an 
acidic  oxide.  Thus  on  the  former  view  the  components  of  copper 
sulphate  are  Cu  and  SO4;  on  the  latter  theory  CuO  and  SO3.  The 
evidence  on  which  these  theories  are  based  and  the  objections  to  the 
Berzelius  theory  have  already  been  fully  considered. 

It  should  further  be  noted  that  the  ionisation  theory  is  only  applic- 
able to  electrolytes  and  is  not  concerned  with  the  great  majority  of 
organic  compounds. 

General  Methods  of  Preparing  Salts— A  number  of 
methods  of  preparing  salts  have  already  been  described  incidentally 
in  the  foregoing  chapters.  The  more  important  general  methods  are 
as  follows : 

(i)  For  Soluble  Salts— Action  of  the  corresponding  acid  (if  a 
moderately  strong  one)  on  the  metal,  oxide,  hydroxide,  or  salt  with  a 
readily  volatile  acid.  Examples  : 

Zn  +  H2SO4->ZnSO4  +  H2  f  , 
KOH  +  HC1-»KC1  +  H2O, 


OXIDES,    ACIDS,    BASES,    AND    SALTS  281 

(2)  For  Insoluble  Salts  —  Double   decomposition   between   a  salt 
containing  the  basic  and  one  containing  the  acidic  constituent,  e.g.  : 

BaCl2+Na2SO4->BaSO4>Ir  +2NaCl. 

(3)  Direct  combination  of  the  elements  (chiefly  used  for  halides 
and  sulphides)  : 


Fe  +  S->FeS. 

(4)  Combination  of  acidic  and  basic  oxides  : 

CaO  +  CO2->CaCO3, 
CaO  +  SiO2->CaSiO3. 

These  methods  are  discussed  more  fully  at  a  later  stage. 

Classification  of  Salts  —  The  usual  definition  of  a  salt  is  a 
substance  formed  by  the  displacement  of  part  or  all  of  the  displaceable 
hydrogen  of  an  acid  by  a  metal.  To  this  it  should  be  added  that 
all  salts  conduct  the  electric  current  in  solution  with  simultaneous 
decomposition  ;  in  other  words,  they  are  electrolytes.  No  subtance 
which  is  not  an  electrolyte  can  rightly  be  termed  a  salt.  Some 
substances  are  rapidly  decomposed  by  water  with  formation  of  sub- 
stances which  conduct  the  electric  current,  and  it  is  very  difficult  to 
determine  whether  the  original  compounds  are  or  are  not  electrolytes. 

We  are  already  familiar  with  two  classes  of  salts  : 

(a)  Normal  Salts,  formed  by  the  displacement  of  all  the  displace- 
able hydrogen  of  an  acid  by  a  metal,  e.g.  Na2SO4  ;  Na3PO4. 

(b)  Acid  Salts,  in  which  only  part  of  the  displaceable  hydrogen  has 
been  replaced  by  a  metal. 

To  this  must  be  added  a  third  class  of  salt,  which  will  now  be 
considered. 

The  formula  of  bismuth  hydroxide,  a  triacidic  base,  is  represented 

/OH  /Cl 

as  follows  :  Bi—  OH,  and  that  of  bismuth  trichloride,  Bi  —  Cl.      The 

\OH  \C1 

salt  can  be  obtained  by  the  action  of  hydrochloric  acid  on  the  base, 
and  may  be  looked  upon  as  being  derived  from  the  base  by  the 
substitution  of  three  Cl  atoms  for  three  OH  groups.  It  now  occurs 

/OH 
to  us  that  compounds  such  as  Bi  —  OH  may  exist,  in  which  only  part 

\C1 

of  the  OH  groups  have  been  displaced  by  the  acid  groups  ;  in  other 
words,  their  composition  is  intermediate  between  that  of  a  normal  salt 


282     A    TEXT-BOOK   OF    INORGANIC   CHEMISTRY 

and  of  a  base.  Such  compounds  are  well  known  ;  they  are  termed 
basic  salts.  Such  a  compound  as  Bi(OH)2Cl  easily  loses  a  molecule  of 
water,  and  the  compound  BiOCl  is  obtained  ;  this  is  also  a  basic  salt. 

Owing  to  the  fact  that  most  basic  salts  are  insoluble  in  water  and 
other  solvents,  their  molecular  weights  cannot  be  determined,  and 
there  is  thus  considerable  uncertainty  as  regards  their  constitutions. 
The  empirical  formula  of  the  basic  chloride  of  bismuth  just  referred 
to  is  BiOCl,  but  it  may  also  be  represented  as  Bi2O3,  BiCl3,  being 
formed  by  the  combination  of  one  molecule  of  the  oxide  and  two 
molecules  of  the  normal  salt.  The  latter  method  is  in  some  respects 
the  most  convenient  way  to  regard  basic  salts.  The  third  type  of  salt 
may  therefore  be  defined  as  follows  : 

(3)  Basic  SaltS)  which  may  be  regarded  as  a  combination  of  one  or 
more  molecules  of  the  normal  salt  with  one  or  more  molecules  of  the 
basic  oxide  or  base,  e.g.  2PbCO3,  Pb(OH)2  (p.  506) ;  BiCl3,  Bi2O3. 

Double  Salts.  Complex  Salts— Reference  has  already  been 
made  to  the  existence  of  so-called  double  salts,  an  excellent  example 
being  carnallite,  MgCl2,KCl,6H2O,  met  with  in  the  Stassfurt  deposits. 
When  carnallite  is  dissolved  in  water,  the  solution  behaves  in  all 
respects  like  a  mixture  of  potassium  and  magnesium  chlorides,  and 
therefore  the  two  salts  exist  separately  in  the  solution.  The  term 
double  salt  is  confined  to  compounds  of  two  or  more  salts  which  exist 
as  such  only  in  the  solid  state,  and  are  completely  broken  up  into 
their  components  on  dissolution  in  water.  In  other  words,  the 
aqueous  solution  of  a  "  double  salt "  contains  only  the  ions  of  the 
simple  salt. 

A  clear  distinction  must  be  drawn  between  double  salts  and  another 
type  of  salts  which  will  now  be  considered.  The  solution  obtained 
by  adding  .potassium  cyanide  to  a  solution  of  silver  nitrate  till  the 
precipitate  first  formed  is  redissolved  clearly  cannot  be  a  mixture  of 
silver  cyanide  and  potassium  nitrate,  because  the  former  compound 
is  insoluble  in  water.  Moreover,  when  an  electric  current  is  passed 
through  the  solution,  the  silver  travels  towards  the  anode,  whilst  the 
potassium,  as  usual,  migrates  towards  the  cathode.  These  results 
show  that  the  silver  forms  part  of  a  negatively  charged  ion,  and 
further  investigation  shows  that  the  ions  are  K'  and  Ag(CN)'2,  so  that 
a  new  substance  is  present,  with  properties  entirely  different  from 
those  of  the  components.  Double  compounds  which  in  solution  give 
rise  to  ions  different  from  those  of  the  components  are  known  as 
complex  salts ,  and  an  ion  resulting  from  the  association  of  a  simple  ion 
with  one  or  more  non-ionised  molecules  is  known  as  a  complex  ion. 


OXIDES,    ACIDS,   BASES,   AND   SALTS  283 

A  clear  distinction  must  be  drawn  between  double  salts,  such  as 
carnallite,  MgCl2,KCl,6H2O,  and  solid  solutions  or  isomorphous  mix- 
tures,  such  as  those  obtained  by  crystallizing  together  a  mixture  of 
sodium  phosphate  and  arsenate  (p.  119).  In  double  salts  the  com- 
ponents are  present  in  simple  molecular  proportions,  whereas  the 
composition  of  a  solid  solution  varies  within  certain,  in  some  cases 
very  wide,  limits  (p.  199). 

Equivalents  of  Acids  and  Bases— The  process  of  neutraliza- 
tion of  acids  by  bases  has  already  been  considered  on  more  than  one 
occasion,  and  it  has  been  pointed  out  that  the  point  at  which  an 
acid  has  been  completely  neutralized  by  a  base,  with  formation  of  a 
neutral  salt,  can  be  recognized  by  adding  a  little  litmus  to  the  solution. 
Substances  of  this  type,,  which  indicate  the  completion  of  a  chemical 
change,  are  known  as  indicators.  The  indicator  to  be  used  depends 
on  the  type  of  chemical  change  which  is  taking  place.  For  the 
neutralization  of  acids  by  bases  three  indicators  are  in  common 
use — (a)  litmus,  which  is  turned  red  by  acids  and  blue  by  alkalis ; 
(b)  phenolphthalein,  which  is  colourless  in  acid,  deep  red  in  alkaline 
solution ;  (c)  methyl  orange,  which  is  yellow  in  alkaline,  red  in  acid 
solution. 

The  quantitative  investigation  of  the  neutralization  of  a  series  of 
acids  by  different  bases  has  led  to  very  interesting  results  Suppose 
we  determine  the  relative  amounts  of  sodium  and  potassium  hydroxide 
required  to  neutralize  a  fixed  amount  of  hydrochloric  acid  and 
find  the  ratio  x  \y.  If  now  we  determine  the  relative  amounts  of 
the  same  bases  required  to  neutralize  a  fixed  amount  of  nitric  acid 
we  find  exactly  the  same  ratio.  Further  investigation  shows  that  we 
have  here  a  general  rule — the  relative  combining  power  of  bases  for 
acids  is  characteristic  for  each  base  and  independent  of  the  nature  of 
the  acid  by  which  it  is  measured,  and  conversely  for  acids  and  bases. 
A  little  consideration  shows  that  we  have  here  another  example  of 
the  law  of  combining  weights  or  equivalents,  which  is  valid  for  com- 
pounds just  as  for  elements.  Each  base  has  a  definite  equivalent, 
which  is  a  measure  of  its  neutralizing  power  for  acids,  and,  similarly, 
each  acid  has  a  definite  eqtiivalent  which  measures  its  combining 
power  for  bases.  This  law  was  discovered  by  Richter  as  early  as 
1791  ;  he  was  Jed  to  it  by  the  observation  that  on  mixing  solutions  of 
two  neutral  salts  no  change  in  neutrality  occurs  even  when  double 
decomposition  takes  place. 

It  only  remains  to  choose  a  unit  to  which  the  combining  weights  of 
acids  and  bases  can  be  referred.  The  most  convenient  for  our  pur- 


284     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

pose  is  to  take  as  the  equivalent  of  an  acid  that  amount  which  con- 
tains i  gram  of  displaceable  hydrogen,  and  as  the  equivalent  of  a 
•base  that  amount  which  neutralizes  one  equivalent  of  an  acid.  The 
meaning  of  this  will  be  clear  from  a  consideration  of  the  equations 
representing  the  neutralization  of  sodium  hydroxide  by  hydrochloric 
and  sulphuric  acid  respectively. 

NaCl  +H2O. 


The  equivalent  of  hydrochloric  acid,  the  'amount  which  contains 
i  gram  of  hydrogen,  is  clearly  36.5  ;  that  of  sodium  hydroxide  is  40. 
As  the  molecular  weight  of  sulphuric  acid  contains  2  grams  of  dis- 
placeable hydrogen,  the  equivalent  of  sulphuric  acid  is  half  the 
molecular  weight  49.  It  is  evident  from  the  equation  that  this  is  in 
accord  with  Richter's  law,  as  those  are  the  relative  amounts  of  the 
two  acids  which  neutralize  the  same  amount  of  sodium  hydroxide. 

Earlier  in  the  chapter  it  has  been  mentioned  that  the  relative 
strengths  of  two  acids  can  be  determined  by  finding  the  ratio  in 
which  an  amount  of  base  insufficient  to  saturate  both  of  them  dis- 
tributes itself  between  them.  Another  method  of  comparing  their 
strengths  would  be  to  measure  their  effect  on  the  rate  of  hydrolysis  of 
cane-sugar  in  aqueous  solution.1  Contrary  to  what  might  at  first 
sight  be  supposed,  their  activities  cannot  be  estimated  by  rinding  the 
relative  quantities  of  a  base  required  to  saturate  definite  quantities  of 
the  acids.  As  a  matter  of  'fact  ',  equal  amounts  of  normal  solutions  of 
both  strong  and  weak  acids  (provided  the  latter  are  not  too  weak} 
require  exactly  the  same  quantity  of  a  strong  base  for  neutralization 
when  a  suitable  indicator  is  chosen.  The  explanation  of  this  fact  in 
terms  of  the  ionisation  theory  will  be  most  readily  understood  from  a 
concrete  case.  We  have  already  seen  that  hydrochloric  acid  is 
almost  completely  ionised  in  dilute  solution,  and  that  the  neutraliza- 
tion by  sodium  hydroxide  consists  essentially  in  the  combination  of 
H*  and  OH'  to  form  water.  Just  as  certain  actions  proceed  in  a 
particular  direction  owing  to  the  formation  of  a  volatile  or  an  in- 
soluble substance,  so  the  driving  force  of  neutralization  is  the 
tendency  of  H'  and  OH'  to  combine  almost  completely  to  form 

1  For  a  full  discussion  of  this  subject  see  Physical  Chemistry,  p.  269. 


OXIDES,   ACIDS,    BASES,    AND   SALTS  285 

non-ionised  water.  Consider  now  the  neutralization  of  a  weak  acid, 
such  as  hypochlorous  acid,  by  sodium  hydroxide,  represented  by  the 
following  equations  — 

(1)  HCIO^H'  +  CIO'. 

(2)  H'+ClO'  +  Na-  +  OH'$Na-  +  ClO/  +  H2O. 

In  hypochlorous  acid  the  H'  concentration  is  comparatively  small, 
but  is  reduced  practically  to  zero  by  the  addition  of  OH'  (as  sodium 
hydroxide).  More  of  the  non-ionised  acid  immediately  dissociates,  the 
fresh  H*  is  also  neutralized  by  OH',  and  the  reaction  proceeds  in  this 
way,  as  represented  by  the  upper  arrows  in  equations  (i)  and  (2),  till 
neutralization  is  practically  complete.  When  a  drop  or  two  of  the 
sodium  hydroxide  has  been  added  in  excess,  the  OH'  ions  cause  a 
change  in  the  colour  of  the  indicator,  indicating  the  completion  of  the 
reaction. 

When,  however,  an  extremely  weak  acid,  such  as  hydrocyanic  acid, 
is  used,  neutralization  is  not  complete  when  an  equivalent  of  alkali 
has  been  added,  owing  to  the  decomposing  influence  of  water  on  the 
salt.  The  equation  under  these  conditions  is  as  follows  : 


or,  in  terms  of  the  ionic  hypothesis  — 


Since  hydrocyanic  acid  is  very  slightly  ionised  and  sodium  hydroxide 
strongly  ionised  the  solution  contains  excess  of  OH'  ions  and  will 
therefore  react  alkaline  to  an  indicator  when  an  equivalent  of  the 
alkaline  has  been  added.  It  will  be  evident  from  these  considerations 
that  an  extremely  weak  acid  cannot  be  estimated  satisfactorily  by 
titration  with  an  alkali,  and,  conversely,  an  extremely  weak  base 
cannot  be  accurately  determined  by  means  of  volumetric  analysis. 

The  above  considerations  form  a  further  illustration  of  the  fact  that 
hydrolysis  may  be  regarded  as  incomplete  neutralization.  Hydrolysis 
has  hitherto  (p.  267)  been  ascribed  to  the  decomposing  influence  of 
water  on  the  salt  in  question.  An  alternative  method  of  regarding  the 
matter  is  that  incomplete  neutralization  in  the  above  case  is  due  to 
the  small  proportion  of  H'  given  by  hydrocyanic  acid,  so  that  a  fairly 
high  concentration  of  the  non-ionised  acid  may  exist  side  by  side  with 
a  moderately  high  concentration  of  OH'.  The  two  explanations  are 
in  essence  the  same,  since  in  both  cases  it  is  a  question  of  the  strength 
of  the  acid. 


286     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Volumetric  Analysis.  Normal  Solutions—  It  is  a  familiar 
fact  that  a  series  of  measurements  of  volume  can  be  carried  out  much 
more  rapidly  and  conveniently  than  an  equal  number  of  weighings.  On 
this  principle  depends  the  utility  of  volumetric  analysis,  which  consists 
in  measuring  the  volume  of  a  solution  of  known  strength  which  is 
required  to  complete  a  definite  chemical  change.  From  the  results 
the  amount  of  the  other  substance  participating  in  the  chemical 
change  can  be  calculated  as  described  below.  The  completion  of  the 
reaction  is  judged  by  means  of  a  suitable  indicator. 

The  solution  of  known  strength  is  termed  a  standard  solution.  It 
may  be  of  any  convenient  strength,  but  is  often  made  to  contain  an 
equivalent  or  a  simple  fraction  of  an  equivalent  per  litre.  A  solution 
which  contains  a  gram  equivalent  of  the  solute  in  a  litre  is  termed  a 
normal  solution.  One  containing  ^  equivalent  is  termed  a  deci- 
normal  solution,  one  containing  T^y  equivalent  a  centinormal  solu- 
tion, and  so  on.  Thus  a  normal  solution  of  hydrochloric  acid  contains 
36.5  grams,  a  decinormal  solution  3.65  grams,  and  a  centinormal 
solution  0.365  grams  of  the  acid  per  litre.  A  decinormal  solution 

of  barium  hydroxide,  Ba(OH)2,  contains    -•      —  =8.5   grams  of  the 

hydroxide  per  litre. 

The  method  of  calculating  the  results  will  be  clear  from  an  example. 
Suppose  25.5  c.c.  of  a  solution  of  barium  hydroxide  containing  10 
grams  of  the  base  per  litre  are  required  to  neutralize  20  c.c.  of  a 
solution  of  nitric  acid  of  unknown  strength.  The  equation  is  as 
follows  : 


171  126 

A  normal  solution  of  Ba(OH)2  contains  85.5  grams  of  the  hydroxide  per  litre 

.  '  .  the  given  solution  is  g  —  normal. 
Since  25.5  c.c.  of  Ba(OH)2  neutralize  20  c.c.  of  the  HNO3  solution,  the  equivalent 

strength  of  the  latter  is  the  greater  in  the  ratio  2-^ 

20 

and  therefore  the  nitric  acid  solution  is  —  X25'5=0  UQ  N 

85.5      20 

Now  a  normal  solution  of  nitric  acid  contains  63  grams  per  litre 

.  '.  the  given  solution  of  the  acid  contains  63  X  o.  149=9.4  grams  per  litre. 

Only  certain  types  of  reaction  are  suitable  for  volumetric  measure- 
ment.    Besides  the  neutralization  of  acids  and  alkalis,  certain  oxida- 


OXIDES,   ACIDS,    BASES,   AND    SALTS  287 

tion-reduction  reactions  and  so-called  precipitation  reactions  in  which 
an  insoluble  substance  is  formed  are  among  those  adapted  for 
volumetric  analysis.  For  details  books  on  volumetric  analysis  should 
be  consulted. 

It  remains  to  consider  the  term  normal  solution  as  applied  to  an 
oxidizing  agent.  Taking  potassium  permanganate,  KMnO4,  as  an 
example,  we  have  seen  (p.  141)  that  in  acid  solution  2 KMnO4  yields  5 
atoms  of  oxygen  for  oxidizing  purposes.  Since  50  is  equivalent  to 
loH,  the  quantity  represented  by  the  formula  2KMnO4 — 316  grams — 
corresponds  to  10  equivalents  of  hydrogen,  and  therefore  by  definition 
a  normal  solution  of  potassium  permanganate  contains  316/10  = 
31.6  grams,  and  a  decinormal  solution  3.16  grams  of  the  salt.  The 
same  considerations  apply  to  other  oxidizing  agents  and  to  reducing 
agents. 

Gravimetric  Analysis.  Determination  of  Atomic 
Weights — Volumetric  analysis,  discussed  in  the  last  section,  is  a 
branch  of  quantitative  analysis,  the  primary  object  of  which  is  the 
determination  of  the  weights  of  the  elements  in  a  given  compound.  In 
volumetric  analysis  the  determination  of  the  quantity  of  the  element  or 
group  is  based  upon  measurements  of  volume.  In  gravimetric  analysis, 
the  other  important  branch  of  quantitative  analysis,  the  proportion  of 
the  element  or  group  in  a  given  compound  is  determined  by  weight. 

For  the  determination  of  the  proportion  of  an  element  in  a  given 
compound  by  weight  two  methods  might  be  suggested,  (i)  the  element 
might  be  isolated  as  such  from  a  definite  quantity  of  the  compound 
and  weighed  ;  (2)  the  element  might  be  isolated  in  the  form  of  a  com- 
pound the  composition  of  which  is  known  and  its  weight  found  by  cal- 
culation. In  practice  the  first  method  is  very  rarely  applicable,  and  the 
second  method  is  almost  invariably  employed.  If,  as  usual,  the  com- 
pound has  to  be  obtained  by  precipitation,  the  following  conditions  must 
be  satisfied  in  accurate  work  :  (i)  the  precipitate  must  be  insoluble  or 
very  slightly  soluble,  otherwise  precipitation  will  be  incomplete  and 
there  will  be  loss  of  weight  on  washing  ;  (2)  the  liquid  used  for  washing 
the  precipitate  must  not  dissolve  it  or  exert  any  chemical  action  on  it ; 
(3)  the  precipitate  has  finally  to  be  dried  and  ignited  to  drive  off  the 
liquid  used  for  washing,  and  must  therefore  be  able  to  withstand  a 
fairly  high  temperature  without  change. 

As  an  illustration  of  the  calculation  of  the  results  we  may  take  the 
determination  of  the  amount  of  barium  in  a  compound  by  precipita- 
tion as  the  sulphate.  Suppose  x  grams  of  the  original  substance  after 
precipitation  with  excess  of  sulphuric  acid  yields  y  grams  of  washed 


288     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

and  dried  barium  sulphate,  BaSO4.  Now  it  can  be  calculated  from 
the  atomic  weights  that  233.44  grams  of  barium  sulphate  contain 
137-37  grams  of  barium,  hence  the  amount  of  barium  in  y  grams  of 


the  sulphate  is  y  x       'f   ^^h   js  tjje   quantity  of  barium    in  x 

grams  of  the  original  substance.  If  the  barium  is  present  in  the 
original  substance  in  the  form  of  a  salt  of  known  composition,  the 
proportion  of  the  latter  in  the  original  substance  can  readily  be 
calculated. 

The  most  accurate  gravimetric  analyses  are  carried  out  in  connection  with 
atomic  weight  determinations,  and  in  this  connection  a  brief  account  of  a  recent 
classical  investigation  by  Richards  and  Willard1  will  be  given,  as  it  shows  also 
how  certain  of  the  more  important  atomic  weights  have  been  determined  with 
reference  to  the  standard  O=i6.  The  main  object  of  the  investigation  was  the 
determination  of  the  ratio  Ag  :  O  through  lithium  chloride  and  perchlorate. 

In  the  first  part  of  the  investigation  the  ratios  of  lithium  chloride  to  silver 
chloride  and  to  silver  were  determined.  A  weighed  amount  of  the  pure  dry 
lithium  chloride  was  dissolved  in  water,  silver  nitrate  added,  and  the  precipitated 
silver  chloride  washed,  dried,  and  weighed.  For  the  description  of  the  numerous 
corrections  the  original  paper  should  be  consulted.  The  results  of  two  experi- 
ments are  given  : 

At.Wt.  of  Li 

LiQ  (grams).         AgCl  (grams).          LiCl/AgCl.         [Ag=  107.88]. 
6.28662  21.25442  0.295779  6.9391 

5.82076  19-67873  0.295790  6.9407 

From  the  results  of  seven  such  experiments  the  value  Li  =  6.g4oi  is  obtained,  with 
a  mean  error  of  +  0.0006  and  a  probable  error  of  +_  0.0002. 

The  ratio  LiCl  :  Ag  was  then  determined  by  adding  silver,  dissolved  in  nitric 
acid,  to  a  solution  of  lithium  chloride  till  neither  Cl'  nor  Ag'  were  in  excess  in  the 
supernatant  liquid.  The  results  of  two  typical  experiments  are  as  follows  : 

At.  Wt.  of  Li 

LiCl  (grams).         Ag  (grams).  LiCl/Ag.  [Ag=  107.88]. 

5.82422  14.82035  0.392988  6.9386 

6.28662  15.99687  0.392991  6.9389 

The  mean  of  seven  experiments  by  this  method  gives  Li  =6.9390,  with  a  probable 
error  of  less  than  +  0.0002. 

In  the  second  part  of  the  investigation,  the  ratio  LiCl  :  40  was  determined  by 
conversion  of  known  weights  of  lithium  chloride  to  perchlorate  by  evaporating 
with  a  slight  excess  of  perchloric  acid.  The  results  of  two  typical  experiments 
are  appended  : 

LiCl  (grams).          LiClO4  (grams).        4O/LiCl.       4O/Ag.        At.  Wt.  of  Silver. 
5.09744  12.79265  1.50962       0.593276  107.876 

4.20534  10.55416  1.50970       0.593307  187.870 

The  numbers  in  column  4  are  obtained  by  multiplying  those  in  column  3  by  the 
ratio  LiCl  :  Ag=o.  392997  :  i,  and  those  in  column  5  by  dividing  40=64  by  the 
numbers  in  column  4.  The  mean  value  is  Ag=  107.  871  +  0.003. 

The  above  results  illustrate  the  extraordinary  accuracy  with  which  gravimetric 
determinations  can  be  carried  out  when  all  possible  precautions  are  taken. 

1  Jour.  Amer.  Chem.  Soc.,  1910,  vol.  xxxii.  p.  4. 


CHAPTER   XXI 

SULPHUR,  SELENIUM  AND  TELLURIUM 
SULPHUR 

Symbol,  S.          Atomic  weight=32.  Molecular  weight=64  (at  1000°). 

/Chemical  Relations — The  compounds  of  sulphur  are  in  many 
^-^  cases  of  the  same  type  as  those  of  oxygen,  mainly  because  both 
elements  are  divalent  Sulphur,  however,  is  also  quadrivalent  (as 
indeed  oxygen  is  in  some  compounds)  and  sexavalent.  The  best 
known  hydrogen  compound  of  sulphur  is  H2S  (corresponding  with 
H2O),  which  is  a  weak  acid.  The  best  known  oxides  have  the  respec- 
tive formulae  SO2,  sulphur  dioxide,  and  SO3,  sulphur  trioxide  ;  the 
acid  corresponding  with  the  former  is  H2SO3,  sulphurous  acid,  whilst 
the  most  important  acid  derived  from  SO3  is  sulphuric  acid,  H2SO4. 

Occurrence — Sulphur  has  been  known  from  the  earliest  times, 
as  it  occurs  free  in  nature,  more  particularly  in  volcanic  districts  in 
Italy,  Sicily,  Iceland,  China,  India  and  in  the  Yellowstone  Park  in 
America.  It  is  probably  formed  by  interaction  of  hydrogen  sulphide 
and  sulphur  dioxide : 

2H2S  +  SO2->2H2O  +  38  >Jr 

both  of  these  gases  being  given  off  from  active  volcanoes.  It  also 
occurs  as  sedimentary  deposits,  not  formed  by  volcanic  action  but 
partly  at  least  by  the  agency  of  reducing  bacteria.  These  deposits 
are  particularly  extensive  in  Texas  and  Louisiana,  and  are  now 
the  source  of  a  considerable  proportion  of  the  world's  sulphur 
supply. 

In  combination  with  hydrogen,  sulphur  occurs  as  hydrogen  sulphide 
in  certain  mineral  springs.  It  also  occurs  very  largely  in  combina- 
tion with  metals,  forming  sulphides.  Some  important  sulphides  are 
galena,  PbS,  iron  pyrites,  FeS2,  copper  pyrites,  CuFeS2,  zinc  blende, 
ZnS,  stibnite  or  antimony  sulphide,  Sb2S3,  and  cinnabar,  which  is 
mercuric  sulphide,  HgS. 

It  also  occurs  as  sulphates,  the  more  important  being  gypsum  or 

19  28g 


29o    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

calcium  sulphate,  CaSO4,  heavy  spar  or  barium  sulphate,  BaSO4,  and 
magnesium  sulphate  in  the  form  of  kieserite  and  Epsom  salts. 

Sulphur  is  one  of  the  essential  constituents  of  albumen,  and  is  there- 
fore found  in  all  plant  and  animal  tissues. 

Preparation — (i)  On  the  commercial  scale,  the  sulphur  is 
separated  from  the  gypsum,  limestone  and  other  substances  with 
which  it  is  usually  mixed,  by  heating  out  of  contact  with  air,  when  it 
melts  and  flows  away  from  the  impurities.  In  Sicily  the  fusion 


FIG.  55. 

is  effected  by  the  wasteful  method  of  burning  a  portion  of  the  sulphur. 
The  sulphur  thus  obtained  is  further  purified  by  distillation,  as  shown 
in  Fig.  55.  It  is  first  melted  in  the  iron  vessel  A,  and  drawn  off  as 
required  into  B,  in  which  it  is  heated  to  boiling,  and  the  vapour  con- 
veyed into  the  large  bricklined  chamber  C.  The  first  portions  of 
vapour  entering  the  chamber  condense  on  the  walls  in  the  form  of  a 
fine  powder,  known  as  flowers  of  sulphur.  As  the  walls  become  hot, 
however,  the  sulphur  melts  and  collects  on  the  floor  of  the  chamber 
as  an  amber-coloured  liquid,  which  is  drawn  off  from  time  to  time  at 


SULPHUR,  SELENIUM  AND  TELLURIUM       291 

D  and  cast  into  sticks  in  wooden  moulds.  In  this  form  it  is  known 
as  roll  sulphur. 

In  Louisiana,  where  sulphur  occurs  400  feet  below  the  surface, 
it  is  melted  by  superheated  steam  and  forced  up  through  a  pipe. 

(2)  By  heating  iron  pyrites  out  of  contact  with  air  and  condensing 
the  vapour  : 


(3)  Sulphur  is  also  obtained  commercially  from  the  "  alkali  waste  " 
of  the  Leblanc  soda  process  (p.  403),  which  contains  a  large  propor- 
tion of  calcium  sulphide,  CaS  ;  and  also  from  the  oxide  of  iron  used 
in  removing  sulphur  compounds  from  coal  gas,  and  which  finaUy  con- 
tains a  considerable  proportion  of  sulphur.  These  methods  are  briefly 
described  at  a  later  stage  (pp.  349  and  405). 

Physical  Properties  —  Sulphur  usually  occurs  as  a  pale  yellow, 
brittle,  crystalline  solid,  insoluble  in  water,  readily  soluble  in  carbon 
disulphide,  and  more  or  less  soluble  in  certain  organic  liquids  such  as 
turpentine,  benzene  and  chloroform.  It  is  a  very  bad  conductor  of 
heat  and  electricity.  This  ordinary  form  melts  at  113°  to  a  pale 
yellow  mobile  liquid.  When  further  heated  to  160°  the  liquid  darkens, 
and  is  so  viscous  that  the  vessel  containing  it  may  be  inverted  without 
causing  it  to  run  out.  At  260°  the  liquid  is  distinctly  less  viscous,  at 
400°  it  is  quite  mobile,  at  445°  it  boils,  giving  off  a  yellowish-brown 
vapour. 

Just  above  its  boiling-point  and  under  fairly  high  pressure  the 
density  of  the  vapour  approximates  to  128,  corresponding  with  the 
formula  S8,  from  860°  to  at  least  1600°  it  is  32,  corresponding  with  the 
formula  S2,  and  it  has  quite  recently  been  shown  that  in  the  neigh- 
bourhood of  2000°  it  falls  to  24  (Nernst),  indicating  that  the  diatomic 
molecules  are  then  split  up  to  some  extent  into  single  atoms.  In  solu- 
tion in  carbon  disulphide  the  formula  of  sulphur  is  S8. 

Sulphur  exists  in  at  least  four  allotropic  modifications,  two  of  which, 
the  "rhombic"  and  "prismatic"  forms,  are  crystalline  and  another 
amorphous. 

Allotropic  Modifications—  (a)  Rhombic  Sulphur—  This  is 
the  form  which  is  stable  at  the  ordinary  temperature,  and  therefore 
native  sulphur  and  ordinary  roll  sulphur  occur  in  this  form.  Good 
crystals  are  obtained  by  the  slow  evaporation  of  a  solution  in  carbon 
disulphide  at  room  temperature  :  they  are  orthorhombic  pyramids. 
The  density  of  rhombic  sulphur  is  2.06  ;  it  melts,  as  already  mentioned, 
at  113°.  This  form  of  sulphur  is  conveniently  designated  Sr. 

(V)  "Monoclinic"  or  "Prismatic"  Sulphur,  Sa—  When  sulphur  is 
melted  in  a  crucible,  the  mass  allowed  to  cool  till  a  crust  forms, 


292     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

and  the  still  liquid  portion  poured  off  through  a  hole  pierced  in  the 
crust,  the  inside  of  the  crucible  is  found  to  be  lined  with  long  needle- 
shaped  crystals.  The  crystals  are  almost  colourless,  the  density  is 
1.9*6,  melting-point  119",  and  they  belong  to  the  monoclinic  system. 

Monoclinic  sulphur  is  stable  above  96°,  but  at  lower  temperatures 
gradually  changes  to  rhombic  sulphur.  Just  as  water  is  stable  above 
o°  and  ice  below  o°,  while  the  two  remain  in  equilibrium  at  that 
temperature,  so  96°  is  the  transition  temperature^  at  which  rhombic 
and  monoclinic  sulphur  are  in  equilibrium,  whilst  at  higher  tem- 
peratures monoclinic,  at  lower  temperatures  rhombic,  is  the  stable 
form.  Neither  of  the  changes  is  very  rapid,  so  that  it  is  possible 
to  determine  the  melting-point,  of  rhombic  sulphur,  which  lies 
about  120°,  before  appreciable  change  into  the  other  form  takes 
place. 

(c)  Nacreous  Sulphur,  Sfn,  discovered  by  Gernez  (1884),  is  obtained 
in  lustrous,  needle-shaped  crystals  by  heating  sulphur  to  150°,  cooling 
to  98°,  and  starting  crystallization  by  rubbing  the  inside  of  the  tube 
with  a  glass  rod.     It  melts  at  103.4°. 

There  is  at  least  one  other  crystalline  modification  of  sulphur,  S,v, 
which  forms  tabular  crystals.  When  a  substance  exists  in  more  than 
one  crystalline  form  it  is  said  to  be  polymorphous;  when  in  two 
crystalline  forms  only  it  is  said  to  be  dimorphous  (p.  175). 

(d)  Liquid  Sulphur.     Amorphous  Sulphur.     Plastic  Sulphur — 
When  Sj  is  melted  at  a  low  temperature  and  suddenly  cooled  by 
pouring  into  cold  water,  a  crystalline  mass  is  obtained  almost  entirely 
soluble  in  carbon  disulphide.     When  the  viscous  liquid  obtained  by 
raising  liquid  sulphur  to  a  higher  temperature  is  treated  in  the  same 
way,  a  plastic  amorphous  mass  is  obtained— plastic  sulphur — almost 
insoluble  in  carbon  disulphide.     In  order  to  account  for  these  obser- 
vations it  is  assumed  that  liquid  sulphur  consists  of  two  liquid  modi- 
fications, S\  and  S^,  in  equilibrium.     At  low  temperatures  the  liquid 
consists  almost  entirely  of  S\,  which  on  cooling  forms  Sj ;  the  higher 
the  temperature  the  greater  the  proportion  of  S^,  which  on  being 
quickly  cooled  past  the  temperature  at  which  it  can  crystallize  rapidly 
forms   amorphous    sulphur,   insoluble    in   carbon    disulphide.      The 
amorphous  modification  on  standing  slowly  changes  to  Sx. 

(e)  Colloidal  Sulphur  is  obtained  in  solution  by  slowly  adding  a 
concentrated  solution  of  sodium  thiosulphate  to  cooled  concentrated 
sulphuric  acid  (p.  314),  or  by  passing  hydrogen   sulphide  into  an 
aqueous  solution  of  sulphur  dioxide  (p.  289). 

(/)  Milk  of  Sulphur  is  obtained  by  boiling  slaked  lime  with  sulphur 
and  water  for  some  time,  pouring  off  the  clear  solution  of  calcium 


SULPHUR,   SELENIUM,  AND  TELLURIUM     293 

polysulphides  and  adding  to  it  hydrochloric  acid,  when  a  white  pre- 
cipitate of  sulphur  slowly  deposits.  From  its  behaviour  with  carbon 
disulphide,  milk  of  sulphur  appears  to  be  a  mixture  of  Sj  and  amorphous 
sulphur.  Flowers  of  Sulphur  also  appears  to  be  a  mixture  of  St  and 
solid  S^.  There  is  no  trustworthy  evidence  of  the  existence  of  the 
"  soluble  amorphous  sulphur  "  sometimes  mentioned  in  the  literature. 

Chemical  Properties — Sulphur  combines  directly  with  many 
other  elements,  both  metals  and  non-metals.  1 1  burns  in  air  or  oxygen 
to  sulphur  dioxide,  SO2.  It  combines  directly  with  hydrogen  when 
heated,  forming  hydrogen  sulphide,  H2S.  If  some  sulphur  is  placed  in 
the  bottom  of  a  test  tube,  which  is  loosely  filled  up  with  copper  filings, 
and  the  sulphur  then  boiled,  the  copper  catches  fire  in  the  vapour  and 
burns  to  the  sulphide,  CuS.  Finely-divided  iron  and  sulphur  also 
combine  directly  (p.  4),  giving  out  heat  and  light. 

Sulphur  is  used  in  the  manufacture  of  sulphuric  acid  and  of  sulphur 
dioxide,  and  forms  one  of  the  constituents  of  ordinary  gunpowder. 
It  is  also  used  in  medicine. 

Velocity  of  Crystallization.  Supercooling.  The  Amor- 
phous State — When  observations  of  the  crystallization  of  a  sub- 
stance such  as  melted  sulphur  are  made  at  temperatures  more  and 
more  removed  from  the  melting-point,  it  is  found  that  the  rate  of 
formation  of  the  solid  phase  increases  at  first  with  the  degree  of 
cooling,  attains  a  maximum,  and  at  still  lower  temperatures  is  ex- 
tremely slow.  It  follows  that  by  very  rapidly  cooling  the  liquid  form  to 
a  low  temperature  many  substances  are  obtained  in  the  non-crystalline 
(amorphous)  form.  Amorphous  solids  are  now  usually  regarded  as 
supercooled  liquids  which  have  not  attained  the  stable  crystalline  form. 

COMPOUNDS  OF  SULPHUR  AND  HYDROGEN 

By  far  the  most  important  compound  of  hydrogen  and  sulphur  is 
hydrogen  sulphide,  H2S.  A  number  of  other  compounds  of  these 
elements,  represented  by  the  general  formula  H2S;t,  where  n  stands 
for  the  integral  numbers  from  2  to  7,  also  appear  to  exist,  and  the 
following  are  definitely  known,  viz.,  H2S2,  H2S3,  and  H2S6. 

HYDROGEN  SULPHIDE,  H2S 

Occurrence — This  compound  occurs  in  certain  mineral  waters, 
in  springs  at  Harrogate,  for  example.  It  is  formed  in  the  decay  of 
organic  matter  containing  sulphur  in  the  absence  of  air ;  and  is  mainly 
responsible  for  the  characteristic  odour  of  rotten  eggs. 

Preparation— (i)  By  direct  combination  of  the  components  on 
heating.  The  reaction  is  comparatively  slow,  but  is  finally  practically 
complete  at  310°;  in  higher  temperatures  it  is  more  rapid,  but  less 
complete  (see  below). 


294    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

(2)  By  the  action  of  dilute  hydrochloric  or  sulphuric  acid  on  ferrous 
sulphide  at  room  temperature  : 

FeS  +  2HCl->FeCl2+H2St. 

This  is  the  usual  laboratory  method  for  preparing  the  gas,  which 
may  conveniently  be  done  in  a  Kipp's  apparatus.  It  is,  however, 
always  contaminated  with  free  hydrogen,  owing  to  the  presence  of 
uncombined  iron  in  commercial  ferrous  sulphide.  The  reaction  pro- 
ceeds practically  completely  in  the  direction  of  the  arrow,  mainly 
because  the  hydrogen  sulphide  readily  leaves  the  system  as  a  gas, 
but  partly  also  because  it  is  a  very  weak  acid  (p.  187). 

(3)  The  pure  compound  is  obtained  by  heating  antimony  sulphide 
with  concentrated  hydrochloric  acid  : 

Sb2S8+6HCl-»2SbCl8  +  3H2S. 
It  is  dried  by  passing  over  anhydrous  calcium  chloride. 

Physical  Properties—  Hydrogen  sulphide  is  a  colourless  gas 
with  a  characteristic  disagreeable  odour.  The  liquefied  gas  boils  at 
—  60°  and  the  solid  melts  at  —83°.  It  is  moderately  soluble  in  water  ; 
the  "solubility"  being  4.37  at  o°,  3.58  at  10°,  and  2.9  at  20°.  It  is 
extremely  poisonous  ;  one  part  in  800  of  air  proved  fatal  to  a  dog,  and 
one  in  1500  is  said  to  cause  death  to  birds.  The  best  antidote  is  the 
inhalation  of  very  dilute  chlorine. 

Hydrogen  sulphide  dissociates  on  heating,  according  to  the  equation 


As  the  compound  is  strongly  exothermic,  the  dissociation  increases 
with  rise  of  temperature.  It  amounts  to  2.3  per  cent,  at  627°,  31.7  per 
cent,  at  1137°,  and  76.1  per  cent,  at  1727°. 

Chemical  Properties—  Hydrogen  sulphide  burns  in  air  with 
a  bluish  flame,  forming  water  and  sulphur  dioxide  : 


In  a  limited  supply  of  air  combustion  is  incomplete,  and  free  sulphur 
is  formed  : 


and  it  is  not  unlikely  that  part  of  the  free  sulphur  found  in  nature  is 
produced  in  this  way. 

Aqueous  solutions  of  hydrogen  sulphide  are  very  unstable,  owing  to 
gradual  oxidation  by  the  oxygen  of  the  air,  water  being  formed  and 
sulphur  separating  as  a  precipitate. 


SULPHUR,    SELENIUM   AND   TELLURIUM      295 

In  harmony  with  the  comparatively  small  affinity  between  its  com- 
ponents (as  illustrated  by  its  dissociation  on  heating)  hydrogen 
sulphide  is  a  powerful  reducing  agent.  The  halogens  are  reduced  to 
the  corresponding  halhydrogen  acids,  e.g.  —  ' 

2Br2  +  2H2S-»4HBr+S2|, 

and  concentrated  sulphuric  acid  is  reduced  to  sulphur  dioxide  and 

water  : 

H2S  +  H20,S03->2H20  +  S0a+  Si- 

On  this  account  sulphuric  acid  cannot  be  used  to  dry  hydrogen 
sulphide,  but  phosphorus  pentoxide  is  suitable  for  this  purpose. 

Hydrogen  sulphide  in  aqueous  solution  acts  as  a  feeble  acid  towards 
litmus,  and  is,  in  fact,  a  dibasic  acid.  As  with  other  dibasic  acids, 
ionisation  takes  place  in  two  stages  : 


but  the  dissociation  is  small  even  in  the  first  stage  (p.  267),  and  in  the 
second  stage  is  very  minute  indeed. 

With  certain  monacidic  bases,  such  as  the  alkalis,  hydrogen 
sulphide  gives  both  acid  and  normal  salts,  e.g.  KHS  and  K2S.  The 
former  are  prepared  by  passing  excess  of  hydrogen  sulphide  into  a 
solution  of  the  base  : 

KOH  +  H2S->KHS  +  H2O. 

Since  H2S  as  a  monobasic  acid,  although  weak,  is  much  stronger  than 
water,  the  compound  KHS  is  only  slightly  hydrolyzed. 

The  normal  sulphide,  K2S,  is  prepared  by  adding  to  the  acid 
sulphide  an  equivalent  of  KOH,  and  removing  the  water  by  evapora- 
tion : 


When,  however,  the  normal  sulphide  K2S  is  dissolved  in  water,  it  is 
practically  completely  hydrolyzed  according  to  the  equation 


because  the  acid  HS'  is  much  weaker  than  the  water,  and  the  solution 
is,  therefore,  strongly  alkaline. 

While  the  sulphides  of  the  alkali  metals,  and  of  certain  other  metals, 
are  soluble  in  water,  those  of  a  number  of  metals,  such  as  zinc  and 
iron,  are  insoluble  in  water  but  soluble  in  acids,  and  still  another 


±96     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

group  of  metals,  including  copper,  mercury,  and  lead,  yield  sulphides 
insoluble  in  dilute  acids.  The  action  of  hydrogen  sulphide  on  a 
solution  of  zinc  chloride,  for  instance,  may  be  represented  as 
follows  :  — 


but  the  equilibrium  lies  very  near  to  the  left,  and  no  precipitation  of 
zinc  sulphide  occurs.  If,  on  the  other  hand,  an  acidified  solution 
of  copper  chloride  is  used,  the  reaction  is  as  follows  :  — 

CuCl2  +  H2S->CuS^  +  2HC1, 

and  owing  to  the  slight  solubility  of  copper  sulphide  is  practically  com- 
plete in  the  direction  of  the  arrow. 

On  these  facts  is  based  a  method  of  detecting  and  separating  many 
of  the  metals  fully  described  in  books  on  qualitative  analysis.  The 
characteristic  colours  of  many  sulphides  are  also  useful  for  purposes 
of  identification. 

Composition  —  When  a  piece  of  tin  is  heated  in  hydrogen 
sulphide  over  mercury,  the  sulphide  and  free  hydrogen  are  formed, 
and  on  cooling  it  will  be  observed  that  no  change  of  volume  has 
occurred.  It  follows  that  the  molecule  of  hydrogen  sulphide  contains 
two  atoms  of  hydrogen,  and  its  formula  is  H2SX  As  the  molecular 
weight  from  the  results  of  density  determinations  is  34,  and  the 
atomic  weight  of  sulphur  is  32,  the  molecule  of  hydrogen  sulphide 
must  contain  one  atom  of  sulphur,  and  its  formula  is  therefore  H2S. 

HYDROGEN  POLYSULPHIDES 

When  sodium  sulphide  is  heated  for  some  hours  with  varying  proportions  of 
sulphur  in  an  atmosphere  of  hydrogen,  and  the  products  dissolved  in  water,  solu- 
tions are  obtained  containing  the  following  polysulphides  of  sodium  :  Na2S2, 
NajjSg,  Na2S4,  and  Na^Sg.  When  these  solutions  are  allowed  to  flow  into  separate 
quantities  of  dilute  hydrochloric  acid,  cooled  in  a  freezing  mixture,  a  yellow,  oily 
liquid  separates  out.  This  liquid  is  a  mixture  of  persulphides  of  hydrogen,  and 
its  composition  varies  with  that  of  the  sodium  persulphide  used  in  obtaining  it. 
From  the  crude  persulphide,  by  fractional  distillation  under  reduced  pressure, 
Bloch  and  Hb'hn  l  have  recently  obtained  the  disulphide,  H2S2,  and  the  trisulphide, 
H-ySs,  in  a  pure  condition. 

Hydrogen  Bisulphide,  H^,  is  a  pale  yellow,  oily  liquid,  of  density  1.376;  it 
boils  at  74  to  75°.  It  decomposes  slowly  at  the  ordinary  temperature,  rapidly 
on  waiming,  giving  off  hydrogen  sulphide  and  depositing  rhombic  sulphur.  It  is 
slowly  decomposed  by  acids,  very  rapidly  by  alkalis,  and  is  fairly  stable  in  benzene 
solution.  It  is  the  analogue  of  hydrogen  peroxide,  H2O2. 

1  Berichte,  1908,  41,  1961-1985. 


SULPHUR,   SELENIUM   AND   TELLURIUM      297 

Hydrogen  Trisulphide,  H2S3,  is  also  a  pale  yellow,  oily  liquid,  of  density  1.496 
at  15°.  In  chemical  behaviour  it  resembles  the  disulphide,  but  is  less  volatile  and 
more  stable  towards  alkab's. 

Hydrogen  Pentasulphide,  H2S5,  also  doubtless  exists,  but  its  properties  have 
not  been  thoroughly  established.  The  investigation  of  these  sulphides  is  rendered 
very  difficult  owing  to  their  tendency  to  dissolve  sulphur. 

COMPOUNDS  OF  SULPHUR  AND  CHLORINE 

The  following  three  well-defined  compounds  of  sulphur  and  chlorine  are  known, 
with  the  respective  formulae  :  S2C12,  SC12,  SC14. 

Sulphur  Monochloride,  SgCLj,  is  prepared  by  passing  dry  chlorine  over  fused 
sulphur  in  a  retort.  It  is  a  yellow  liquid  which  boils  at  137  to  138°;  the  solid  melts 
at  -75  to  -76°.  It  is  the  most  stable  of  the  chlorides  of  sulphur,  being  only 
slightly  dissociated  at  its  boiling-point.  It  is  readily  decomposed  by  water  : 


It  is  an  excellent  solvent  for  sulphur,  dissolving  over  60  per  cent,  of  the  latter 
at  room  temperature,  and  the  solution  is  used  in  the  vulcanization  of  rubber. 

Sulphur  Bichloride,  SC12,  is  prepared  by  passing  chlorine  into  the  mono- 
chloride,  cooled  to  -  15°,  until  the  theoretical  gain  in  weight  is  attained.  It  is  a 
dark  reddish-brown  liquid  of  density  1.622  at  15°  and  boils  at  59°  under  at- 
mospheric pressure.  At  its  boiling-point  it  is  partially  decomposed  into  the 
monochloride  and  chlorine.  It  is  readily  decomposed  by  water  : 

2SC12+2H20->4HC1+S02+S. 

Sulphur  Tetrachloride,  SC14,  is  obtained  when  the  dichloride  and  chlorine 
are  brought  together  at  temperatures  below  -  75°.  The  pure  compound  is  obtained 
by  freezing  out,  the  excess  of  chlorine  being  removed  by  centrifugal  action.  The 
crystals  melt  between  -  30°  and  -  20°.  Even  at  low  temperatures  it  is  partially 
dissociated  and  decomposition  is  practically  complete  at  room  temperature.  It 
is  readily  decomposed  by  water  : 

SC14+2H2O->SO2+4HCL 

With  bromine,  sulphur  forms  the  compound  S2Br2,  a  brownish-red  liquid, 
which  boils  with  partial  decomposition  at  200°. 

No  compounds  of  sulphur  and  iodine  are  known  with  certainty. 

Sulphur  Hexafluoride,  SF6,  obtained  by  Moissan  by  the  direct  action  of  fluorine 
on  sulphur,  is  a  colourless  gas  which  can  be  condensed  to  a  solid  melting  at  55°  ; 
the  liquid  boils  at  a  slightly  higher  temperature.  The  gas  is  fairly  stable  and 
chemically  inactive.  The  existence  of  this  compound  is  fairly  conclusive  evidence 
that  sulphur  can  act  as  a  sexavalent  element. 

OXIDES  AND  OXYACIDS  OF  SULPHUR 
The  following  four  oxides  of  sulphur  are  known  :  — 

Disulphur  trioxide  (hyposulphurotis  anhydride)  .  .  S2O3 

Sulphur  dioxide  (sulphurous  anhydride)       .        .  .  SO2 

Sulphur  trioxide  (sulphuric  anhydride)         .         .  .  SO3 

Sulphur  heptoxide  (persulphuric  anhydride)        .  .  S2O7. 


298     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

The  following  oxyacids,  the  first  five  of  which  are  derived  from  the 
oxides  just  mentioned,  are  known  : — 

Hyposulpharous  acid        .        .     H2S2O4     Q_£.QTT 

/OH 
Sulphurous  acid        .        .        .     H2SO3       OS\ 

\OH 

/OH 
Sulphuric  acid          .        .        .     H2SO4      O2S<f 

/OH     OH\ 

Pyrosulphuric  acid  .         .         .    H2S2Or     O2S^—    — O SO2 

(Nordhausen  sulphuric  aid) 

/OH 

Permonosulphuric  acid    .         .     H2SO6      O2S/ 

(Caro's  acid)  XO'OH 

/OH    OH\ 

Persulphuric  acid     ,        ,        .     H2S2O8     O2S<  /SO2 

\  O—  O  / 

/OH 

Thiosulphuric  acid  .        .     H2S2O3     OoS< 

\SH. 

Besides  these,  four  less  important  acids,  dithionic  acid,  H2S2Og, 
trithionic  acid,  H2S3O6,  tetrathionic  acid,  H2S4O6,  and  pentathionic 
acid,  H2S6O6,  are  known. 

SULPHUR  DIOXIDE,  SO2 

Occurrence— As  already  mentioned,  sulphur  dioxide  is  given 
off  from  active  volcanoes.  It  also  occurs  in  traces  in  the  atmosphere 
of  towns,  being  formed  from  the  combustion  of  the  sulphur  always 
present  in  coal. 

Preparation — (i)  By  burning  sulphur  in  air  or  oxygen : 


A  little  sulphur  trioxide  (in  air  up  to  7  per  cent.)  is  formed  at  the 
same  time. 

(2)  By  burning  sulphides,  e.g.  iron  pyrites,  FeS2,  in  air  : 

4FeS2+iiO2=2Fe2O3 


SULPHUR,    SELENIUM   AND   TELLURIUM     299 

This  method  is  used  on  the  large  scale  in  the  manufacture  of  sulphuric 
acid. 

(3)  By  heating  copper  turnings  with  concentrated  sulphuric  acid  : 


This  is  the  most  convenient  laboratory  method  for  preparing  the  gas, 
which  can  be  collected  over  mercury  or  by'upward  displacement  of 
air. 

(4)  The  reduction  of  concentrated  sulphuric  acid  to  sulphur  dioxide 
can  also  be  effected  by  heating  it  with  carbon  or  sulphur: 


These  and  the  equations  representing  the  reduction  of  sulphuric  acid 
to  sulphur  dioxide  can  readily  be  obtained  when  the  acid  is  written 
in  the  fofm  SO3,H2O  (cf.  p.  226). 

(5)  By  acting  on  sulphites  with  dilute  sulphuric  or  hydrochloric 
acid  : 


Physical  Properties—  Sulphur  dioxide  is  a  colourless  gas, 
with  the  penetrating  and  characteristic  odour  associated  with  burning 
sulphur.  It  can  easily  be  obtained  as  a  colourless  mobile  liquid,  most 
conveniently  by  passing  the  gas  (obtained  by  the  action  of  sulphuric 
acid  on  copper  and  dried  by  bubbling  through  concentrated  sulphuric 
acid)  through  a  tube  immersed  in  a  freezing  mixture  of  ice  and  salt 
(Fig.  56).  Liquid  sulphur  dioxide  boils  at  -  8°  ;  at  o°  its  vapour 
pressure  is  1.87  atmospheres,  at  20°  it  is  3.24  atmospheres.  It  is 
a  good  solvent,  especially  for  inorganic  salts,  and  many  of  the  solu- 
tions are  good  electrolytes.  On  account  of  its  ready  volatility  and 
large  heat  of  vaporization,  liquid  sulphur  dioxide  is  used  for  obtain- 
ing low  temperatures  by  its  rapid  vaporization  (p.  66).  The  liquid 
is  now  obtainable  commercially  in  syphons.  Solid  sulphur  dioxide 
melts  at  -76°. 

Sulphur  dioxide  is  readily  soluble  in  water,  at  o°  the  "solubility"  is 
79.79,  at  20°  39.37,  and  at  40°  18.766.  It  is  completely  expelled  from 
solution  by  boiling. 

Chemical  Properties  —  The  aqueous  solution  of  sulphur  dioxide 
has  an  acid  reaction,  owing  to  the  fact  that  it  combines  with  water  to 
form  sulphurous  acid,  H2SO3  (see  below). 

Sulphur  dioxide  is  readily  oxidized  to  the  trioxide,  SO3,  and  is 
therefore  a  useful  reducing  agent.  It  combines  directly  with  oxygen 


300     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

at  high  temperatures  (p.  302),  but  its  reducing  character  is  more  pro- 
nounced in  aqueous  solution,  owing  to  the  readiness  with  which 
sulphurous  acid  is  converted  to  sulphuric  acid.  The  free  halogens* 


FIG.  56. 

iodates  (p.  185),  potassium  dichromate  (p.  537),  and  other  oxidizing 
agents  are  readily  reduced  at  the  ordinary  temperature  : 


+  5S02+4H20->5H2S04-fI2. 
Papers  soaked  with  potassium  iodate  and  starch  are  sometimes  used 


SULPHUR,    SELENIUM   AND   TELLURIUM      301 


for  the  detection  of  sulphur  dioxide,  as  they  turn  blue  when  iodine 
is  liberated. 

Sulphur  dioxide  in  the  presence  of  water  is  a  useful  bleaching  agent. 
In  most  cases  this  is  connected  with  its  reducing  properties,  the 
colouring  matters  being  reduced  by  the 
hydrogen  liberated  according  to  the 
equation 

SO2  +  2H2O->H,SO4 


In  other  cases  (e.g.  the  bleaching  of 
flowers)  the  change  appears  to  depend 
upon  direct  combination  of  the  dioxide 
with  the  colouring  matters,  as  the  colour 
is  restored  on  warming  or  on  momen- 
tary exposure  to  the  action  of  chlorine. 
It  is  used  for  bleaching  straw,  wool, 
silk  and  other  materials  that  would  be 
damaged  by  chlorine. 

Sulphur  dioxide,  both  in  the  gaseous 
form  and  in  solution,  is  a  useful  dis- 
infectant. 

Composition  —  The  composition 
of  sulphur  dioxide  can  be  determined 
with  the  apparatus  represented  in 
Fig.  57.  A  small  quantity  of  sulphur  is 
placed  in  the  cup  (connected  to  a 
copper  wire)  which  is  then  placed  in 
position  in  the  bulb,  the  latter  being 
filled  with  oxygen  at  atmospheric 
pressure.  The  position  of  the  mercury 
in  the  right-hand  tube  is  noted,  the 
pressure  reduced  by  running  out  mer- 
cury at  the  stopcock,  and  the  sulphur 
ignited  by  passing  an  electric  current,  L-  -  -  * 
which  heats  to  redness  the  thin  platinum  FIG.  57* 

wire  which  has  been  arranged  to  dip  in  the  sulphur.  After  combustion 
is  complete  and  the  apparatus  has  cooled,  the  pressure  is  again  adjusted 
to  that  of  the  atmosphere  by  pouring  mercury  into  the  left-hand  limb 
and  it  will  then  be  found  that  the  mercury  in  the  other  limb  stands  at 
the  same  level  as  before.  It  follows  that  the  volume  of  sulphur  dioxide 
produced  is  equal  tq  that  of  the  oxygen  used  up  in  its  formation  ; 


3o2     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

in  other  words,  sulphur  dioxide  contains  its  own  volume  of  oxygen. 
By  Avogadro's  hypothesis  the  molecule  of  sulphur  dioxide  therefore 
contains  a  molecule  or  two  atoms  of  oxygen  and  its  formula  is  SxOp 
where  x  is  a  whole  number.  The  molecular  weight  of  the  gas,  deter- 
mined from  its  density,  is  64,  and  as  it  contains  2  atoms  or  32  parts 
of  oxygen,  there  remains  32  parts  of  sulphur.  Now  no  sulphur  com- 
pound is  known  the  molecule  of  which  contains  less  than  32  parts 
of  that  element,  so  that  32  is  the  atomic  weight  of  sulphur,  and  the 
formula  of  sulphur  dioxide  is  SO2. 

Sulphurous  Acid  and  Sulphites—  Sulphurous  acid,  H2SO3, 
has  never  been  obtained  otherwise  than  in  aqueous  solution.  As  a 
dibasic  acid  it  forms  two  series  of  salts  with  alkalis,  the  normal 
sulphites,  for  example  Na2SO3,  and  the  acid  sulphites,  for  example 
NaHSO3,  which  are  soluble  in  water.  The  normal  sulphites  in 
aqueous  solution  are  considerably  hydrolyzed,  as  sulphurous  acid, 
though  a  fairly  strong  monobasic  acid  (HgSOg^tH'  +  HSOgO  is  a  very 
weak  dibasic  acid  (HSO/^tH-  +  SO3"),  though  not  so  weak  as  hydrogen 
sulphide.  The  sulphites  of  the  other  metals  are  mostly  insoluble  in 
water. 

SULPHUR  TRIOXIDE 

Preparation  —  (i)  By  direct  combination  of  sulphur  dioxide  and 
oxygen  : 


The  combination  is  very  slow,  even  at  high  temperatures,  but  is 
fairly  rapid  when  the  mixture  is  passed  over  finely-divided  platinum, 
platinum  asbestos  or  ferric  oxide,  heated  to  400°. 

On  passing  the  vapours  into  a  cooled  receiver,  the  trioxide  is 
obtained  in  colourless  crystals.  At  400°  combination  is  practically 
complete  under  equilibrium  conditions,  but  as  the  temperature  is 
raised  the  trioxide  undergoes  dissociation,  and  the  change  is  practi- 
cally complete  in  the  direction  of  the  upper  arrow  at  1000°.  As  this 
reaction  is  now  employed  in  the  commercial  preparation  of  sulphuric 
acid  it  is  fully  described  at  a  later  stage. 

(2)  By  heating  pyrosulphuric  acid  (fuming  sulphuric  acid)  (q.v.)  : 

H2S207->H2S04  +  S03t. 

(3)  By  strongly  heating  ferric  sulphate  : 


SULPHUR,    SELENIUM   AND   TELLURIUM     303 

(4)  By  dehydrating  concentrated  sulphuric  acid  by  means  of  phos- 
phorus pentoxide  : 


Physical  Properties  —  Sulphur  trioxide  is  a  liquid  at  the 
ordinary  temperature  ;  it  boils  at  46°,  and  gives  dense  white  fumes 
when  exposed  to  the  air,  owing  to  the  combination  of  its  vapour  with 
moisture,  forming  sulphuric  acid.  The  crystals  obtained  by  cooling 
the  liquid  melt  at  14.8°.  A  second,  more  stable  form  of  the  solid 
occurs  in  long,  needle-shaped  crystals  resembling  asbestos.  This 
form  vaporizes  on  heating  without  previously  melting,  and  is  pre- 
sumably a  polymer  (p.  175)  of  the  ordinary  form,  probably  (SO3)2. 
The  vapour  at  low  temperatures  contains  the  two  forms  in  equilibrium, 


but  dissociates  completely  into  the  simpler  form  as  the  temperature 
rises. 

The  dissociation  of  the  trioxide  into  sulphur  dioxide  and  oxygen  at 
high  temperatures  has  already  been  referred  to. 

Chemical  Properties  —  Sulphur  trioxide  unites  very  vigorously 
with  water  to  form  sulphuric  acid  : 

SO3  +  H2O->H2SO4. 

It  also  unites  directly  with  certain  metallic  (basic)  oxides  with  forma- 
tion of  salts  : 

CaO  +  SO3-»CaSO4. 

The  combination  of  basic  and  acidic  oxides  is  one  of  the  general 
methods  for  preparing  salts  (p.  390). 


SULPHURIC  ACID 

History — Sulphuric  acid  was  a  familiar  substance  to  the  al- 
chemists, being  prepared  by  distilling  ferrous  sulphate,  FeSO4,7H2O 
(p.  560).  The  latter  compound  was  known  as  green  vitriol,  whence 
the  name  oil  of  vitriol^  still  sometimes  applied  to  sulphuric  acid. 

Preparation — On  the  commercial  scale  sulphuric  acid  is  pre- 
pared by  bringing  about  the  combination  of  sulphur  dioxide  and 
oxygen,  and  dissolving  the  resulting  sulphur  trioxide  in  water.  The 
difficulty  arising  from  the  extreme  slowness  with  which  the  gases 
combine  directly  (p.  302)  is  overcome  by  using  catalysts.  In  the 


3o4     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

"contact  process,"  which  will  be  first  described,  the  catalyst  is 
platinized  asbestos  ;  in  the  "  lead  chamber  process "  oxides  of 
nitrogen  are  used  for  this  purpose. 

(1)  The  Contact  Process — Although   this  method  is  apparently  a 
very   simple   one,  serious   difficulties  were   encountered  in  practice 
owing  to  the  fact  that  the  platinum  soon  lost  its  catalytic  power. 
This   was   ultimately  traced  to  the  presence  of  impurities — mainly 
arsenic  and  dust — in  the  mixture  of  sulphur  dioxide  and  air  used  in  the 
process.     This  mixture  is  usually  prepared  by  roasting  iron  pyrites 
(p.  289),  which  contain  arsenic,  hence  the  presence  of  this  impurity 
in  the  reaction  mixture.     Methods  for  removing  the  arsenic  and  dust 
have  been  worked  out  by  Knietsch  (1899),  to  whom  the  commercial 
success  of  this  method  is  mainly  due.1 

In  outline  the  contact  process  is  as  follows.  The  purified  mixture 
of  sulphur  dioxide  and  air,  of  such  composition  that  the  ratio  SO2  :  O2 
is  approximately  2  :  3,  is  passed  through  tubes  containing  the  platinized 
asbestos  and  kept  at  400  to  450°.  As  a  large  amount  of  heat  is  given 
out  in  the  process  of  combination,  and  a  higher  temperature  is 
disadvantageous,  the  cooling  is  effected  by  passing  fresh  portions  of 
the  reaction  mixture  round  the  outside  of  the  tubes,  the  mixture 
being  thus  warmed  to  the  most  favourable  temperature  before  enter- 
ing the  reaction  tubes.  By  increasing  or  diminishing  the  rate  of 
flow  of  gas,  the  temperature  can  be  regulated  satisfactorily.  The 
issuing  trioxide  is  not  passed  into  water,  as  part  of  it  is  apt  to  escape 
condensation,  but  is  absorbed  in  97  to  98  per  cent,  sulphuric  acid, 
which  is  kept  at  that  strength  by  the  simultaneous  addition  of  water. 

(2)  The  Lead  Chamber  Process — As  already  indicated,  this  method 
consists  essentially  in  the  oxidation  of  sulphur  dioxide  to  sulphuric 
acid  by  free  oxygen  in  the  presence  of  aqueous  vapour,  oxides  of 
nitrogen  being  used  to  accelerate  the  reaction.     A  mixture  of  sulphur 
dioxide,  air,  aqueous  vapour  and  nitrous  fumes  is  allowed  to  interact 
in   large  leaden   chambers   at   room  temperature,  and  ultimately  a 
dilute  solution  of  sulphuric  acid  collects  on  the  floor  of  the  chambers. 

The  simplest  method  of  representing  the  reactions  is  as  follows: — 

(1)  SO2  +  NO2-»SO3+NO 

(2)  S03  +  H20-»H2S04 

(3)  2NO  +  O2-»2NO2. 

1  Other  inventors,  notably  Winkler  in  Freiburg  and  Messel  in  London,  appear 
also  to  have  solved  the  difficulties  of  the  process  before  the  publication  of  Knietsch's 
patent  in  1899,  but  their  methods  were  kept  secret. 


SULPHUR,    SELENIUM   AND    TELLURIUM     305 

As  sulphur  dioxide  and  nitrogen  dioxide  react  much  more  rapidly 
than  the  former  gas  does  with  oxygen,  we  can  easily  understand 
why  sulphuric  acid  is  obtained  so  readily  by  an  indirect  method.  The 
nitric  oxide  formed  in  equation  (i)  is  rapidly  reconverted  to  the 
peroxide,  according  to  equation  (3),  so  that  theoretically  a  small 
amount  of  peroxide  could  convert  an  unlimited  amount  of  sulphur 
trioxide  to  sulphuric  acid.  In  practice,  however,  secondary  reactions 
also  occur  to  some  extent  (for  example,  reduction  to  nitrous  oxide, 
N2O,  which  is  no  longer  able  to  combine  with  oxygen,  p.  234),  result- 
ing in  the  loss  of  activity  ot  part  of  the  oxides  of  nitrogen,  so  that 
they  have  to  be  continuously  supplied  as  the  action  proceeds. 

The  reactions  taking  place  in  the  lead  chambers  are  probably  more 
complicated  than  the  above  simple  scheme  would  indicate,  but  they 
are  not  yet  thoroughly  understood.  According  to  Lunge,  reaction 
proceeds  mainly  in  two  steps  ;  the  first  leads  to  the  formation  of 

/OH 

nitrosyl-sulphuric   acid,    SO2\  ,  according  to  the  equation 

M3-NO 

/OH 

\O-NO 

and  the  second  stage  consists  in  the  rapid  breaking  up  of  nitrosyl- 
sulphuric  acid  by  water,  with  formation  of  sulphuric  acid,  nitric  oxide, 
and  nitrogen  peroxide  : 

/OH  /OH 

2SO2<  +  H2O->2SO2\          -f-NO  +  NOo. 

\0-NO  \OH 

As  the  formula  shows,  nitrosyl-sulphuric  acid  is  derived  from  sul- 
phuric acid  by  the  displacement  of  one  of  the  hydrogen  atoms  by 
the  univalent  -  N  =  O  group.  This  substance  does  not  appear  when 
the  process  is  working  properly,  but  if  the  supply  of  water  is  in- 
sufficient it  separates  on  the  sides  of  the  chambers  in  colourless 
crystals,  the  so-called  "  chamber  crystals."  In  a  later  paper  Lunge 
modified  this  theory  to  some  extent.1 

Raschig,2  on  the  other  hand,  considers  that  nitrous  acid,  HNO2  or  NOOH,  is 

•  /OH 

the  active  catalyst,  and  that  two  unstable  intermediate  compounds,  SO2\ 

\NO 

1  Cf.  Encyc.  Britannica,  nth  Edition,  vol.  xxvi.  p.  67. 

2  Journ.  Soc.  Chem.  Ind.,  1910. 
20 


306     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


nitroso-sulphuric  acid,  and  the  compound  SO2<^    ^O  have  a   transient  exist- 

\OH 

ence  during  the  change.      Raschig's    view    is    summarized    in  the    following 
equations : — 


(i) 
(2) 

(3)  NO(OH)S02-OH 

(4)  2NO+O  +  H2O->2NO-OH. 


The  production  of  sulphuric  acid  by  this  process  can  be  shown 


AIR 


FIG.  58. 


on  the  small  scale  in  the  laboratory.  A  large  wide-mouthed  flask 
(Fig.  58)  is  provided  with  a  cork  carrying  five  tubes,  three  ol 
which  are  used  for  leading  nitrous  fumes,  sulphur  dioxide  and  air 
into  the  flask.  By  means  of  the  fourth  tube  steam  can  be  introduced, 
and  the  fifth  tube,  which  passes  just  through  the  cork,  serves  as  an 
exit  tube.  When  the  four  substances  are  allowed  to  interact,  sulphuric 
acid  is  formed  and  collects  at  the  bottom  of  the  flask.  If  it  is  desired 
to  show  the  formation  of  chamber  crystals,  nitrous  fumes,  air,  sulphur 
dioxide,  and  a»very  little  water  vapour  are  admitted,  and  the  crystals 
form  on  the  sides  and  bottom  of  the  flask.  When  steam  is  passed 
in,  the  crystals  disappear  and  sulphuric  acid  is  formed. 

The  arrangement  of  the  apparatus  used  on  the  commercial  scale  is 


SULPHUR,    SELENIUM   AND  TELLURIUM      307 

shown  in  Fig.  59.     The  sulphur  dioxide,  obtained  by  burning  sulphur, 


or  more  usually  by  roasting  iron  pyrites  or  other  sulphides  in  a  current 
of  air  (the  pyrites  burners  are  shown  at  P  P),  is  first  led  into  a  long  flue 


3o8     A   TEXT-BOOK    OF   INORGANIC    CHEMISTRY 

in  which  the  dust  particles  mechanically  carried  along  settle  out,  and 
the  proper  proportion  of  air  is  added.  The  mixed  gases,  which  are  at 
a  temperature  of  about  300°,  are  then  passed  up  the  Glover  tower, 
G,  which  is  a  high  tower  lined  inside  with  lead  and  filled  with  stones, 
over  which  dilute  sulphuric  acid  containing  dissolved  oxides  of  nitro- 
gen (nitrosyl-sulphuric  acid)  is  trickling.1  The  nitrosyl-sulphuric  acid 
is  decomposed  by  the  hot  furnace  gases,  with  formation  of  sulphuric 
acid,  which  is  collected  at  the  bottom  of  the  tower  (at  A),  and  nitric 
oxide,  which  is  carried  along  by  the  stream  of  gas,  now  at  a  much 
lower  temperature,  into  the  first  of  the  chambers  : 


/OH 
•NO 


2SO2<  +  SO2+2H2O->3H2SO4 

X" 


The  chambers,  which  are  constructed  completely  of  sheet  lead,  often 
have  a  capacity  of  1 50,000  to  200,000  cubic  feet  each,  and  are  arranged 
in  sets  of  three  or  four  connected  together  (two  only  are  shown  in  the 
figure).  As  the  gases  are  slowly  drawn  through  these  chambers, 
water  in  the  form  of  steam  is  injected,  and  the  chemical  reactions 
already  described  (p.  305)  take  place,  resulting  in  the  formation 
of  sulphuric  acid,  which  collects  on  the  floors  of  the  chambers.  The 
gases  escaping  from  the  last  of  the  chambers,  which  consist  mainly  of 
oxides  of  nitrogen  and  a  large  proportion  of  nitrogen  from  the  air 
o'riginally  drawn  into  the  chamber,  are  passed  up  a  second  tower,  the 
so-called  Gay-Lussac  tower,  H,  which  is  filled  with  coke,  over  which 
80  per  cent,  sulphuric  acid  continually  trickles  from  a  reservoir  at  the 
top  of  the  tower.  The  acid  used  for  this  purpose  is  part  of  that  drawn 
off  at  the  bottom  of  the  Glover  tower,  whence  it  is  pumped  up  to  the 
top  of  the  Gay-Lussac  tower.  The  object  of  the  Gay-Lussac  tower 
is  to  prevent  the  escape  of  the  oxides  of  nitrogen  ;  these  are  absorbed 
almost  completely  by  the  concentrated  acid  with  formation  of  nitrosyl- 
sulphuric  acid.  This  acid  is  then  pumped  to  the  top  of  the  Glover 
tower,  where  the  oxides  of  nitrogen  are  removed  trom  it  by  the  hot 
furnace  gases,  and  swept  back  into  the  chambers,  as  already  fully 
explained. 

Owing  to  secondary  reactions  and  unavoidable  losses,  however,  the 
oxides  of  nitrogen  are  gradually  used  up,  and  the  loss  must  be  made 
good.  This  is  effected  by  heating  sodium  nitrate  with  sulphuric  acid 
in  earthenware  pots,  which  are  placed  in  the  path  of  the  furnace  gases 

1  A  mixture  of  acid  from  the  chambers  and  from  the  Gay-Lussac  tower  is  used 
for  this  purpose. 


SULPHUR,    SELENIUM   AND   TELLURIUM      309 

on  their  way  to  the  chamber  at  N  N.     The  nitric  acid  is  rapidly 
reduced  to  nitrogen  peroxide  by  the  sulphur  dioxide  : 


When  the  sulphuric  acid  in  the  chamber  has  reached  a  density  of 
1.  60  (68  per  cent,  of  acid)  it  is  withdrawn  (more  concentrated  acid 
attacks  the  material  of  the  chambers  and  dissolves  oxides  of  nitrogen) 
and  further  concentrated  to  a  density  of  1.70  (77  per  cent,  of  acid) 
by  evaporation  in  leaden  pans.1  Above  this  strength  lead  would 
be  dissolved  fairly  readily,  so  the  final  concentration  is  effected  by 
heating  in  glass,  platinum,  or  cast-iron  vessels.  The  commercial  acid 
thus  obtained  has  a  density  of  1.83  to  1.84,  and  contains  93  to  98  per 
cent,  of  the  pure  acid. 

The  commercial  acid  contains  a  number  of  impurities,  more 
particularly  arsenic  from  the  pyrites,  lead  sulphate,  and  oxides  of 
nitrogen.  Most  of  the  impurities,  except  the  arsenic,  are  got  rid  of 
by  adding  a  little  ammonium  sulphate  and  distilling  from  platinum 
stills  ;  the  arsenic  is  removed  by  means  of  hydrogen  sulphide. 

Physical  Properties  —  The  purest  sulphuric  acid  which  can 
be  obtained  by  distillation  contains  1.5  per  cent,  of  water,  and  its 
density  is  1.842  at  15°.  The  anhydrous  acid  is  most  conveniently 
prepared  by  adding  the  theoretical  amount  of  sulphur  trioxide  to 
the  98.5  per  cent,  acid  ;  it  is  a  colourless,  odourless,  oily  liquid 
of  density  1.838  at  15°.  When  the  pure  acid  is  heated,  it  begins 
to  decompose  about  150°,  giving  oft"  white  fumes  of  sulphur  tri- 
oxide. At  338°,  when  the  concentration  has  fallen  to  98.5  per 
cent,  of  acid  through  loss  of  sulphur  trioxide,  the  liquid  boils. 
The  98.5  per  cent,  acid  is  therefore  a  constant-boiling  mixture, 
similar  to  that  formed  by  hydrochloric  acid  and  water.  Even  at 
its  boiling-point,  the  vapour  of  sulphuric  acid  is  dissociated  to  a 
considerable  extent  into  the  trioxide  and  water  vapour  : 

S03+H2O^H2S04. 

Chemical  Properties  —  When  sulphuric  acid  is  added  to 
water,  a  great  quantity  of  heat  is  evolved.  When  i  mol  of 
water  is  added  to  i  mol  (98  grams)  of  acid,  6700  cal.  are  given 
out.  More  heat  is  given  out,  but  in  diminishing  amount,  as  further 
quantities  of  water  are  successively  added.  The  total  heat  evolved 
when  a  mol  of  acid  is  added  to  a  very  large  excess  of  water  is  about 

1  Chamber  acid  can  also  be  concentrated  up  to  78-80  per  cent,  by  passing  it 
through  the  Glover  tower. 


3io     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

20,000  calories.  This  remarkable  thermal  effect  is  doubtless  due 
mainly  to  chemical  combination  between  the  acid  and  water, 
and  partly,  in  all  probability,  to  progressive  electrolytic  dissociation. 
One  well-defined  hydrate  of  sulphuric  acid,  H2SO4,H2O,  has  been 
obtained  in  prismatic  crystals,  melting  at  8.5°.  There  is  some  evi- 
dence of  the  existence  of  a  second  hydrate,  H2SO4,2H2O,  but  it  has 
not  been  definitely  isolated.  Sulphuric  acid  itself  may  of  course  be 
regarded  as  the  monohydrate  of  sulphur  trioxide.  It  is  interesting 
to  note  that  whilst  in  the  reaction  SO3+  H2O->H2SO4  21,300  calories 
are  given  out,  the  combination  with  a  further  molecule  of  water,  to 
form  SO3,2H2O  or  H2SO4,H2O,  liberates  only  6700  calories. 

Further  evidence  of  the  great  affinity  between  concentrated  sul- 
phuric acid  and  water  is  to  be  found  in  the  fact  that  it  abstracts  the 
elements  of  water  from  many  compounds.  In  the  case  of  organic 
compounds  such  as  sugar  and  paper,  which  are  composed  of  carbon, 
hydrogen,  and  oxygen,  the  removal  of  the  last  two  elements  leads 
to  the  liberation  of  free  carbon,  and  the  substances  are  said  to  be 
charred.  The  use  of  sulphuric  acid  for  drying  gases  has  already 
been  repeatedly  referred  to.  Sulphuric  acid  is  also  largely  used  for 
facilitating  actions  in  which  water  is  one  of  the  products  (e.g.  the 
preparation  of  nitro-glycerine)  ;  it  takes  up  the  water  and  prevents  a 
reverse  reaction  by  which  the  product  required  would  be  decomposed. 

In  aqueous  solution  sulphuric  acid  acts  as  a  dibasic  acid.  As  with 
other  dibasic  acids,  it  ionises  in  two  stages  : 

(i)  H2S04^H'  +  HS04'    (2)  HS04'^H'  +  S04" 

the  second  stage  lagging  considerably  behind  the  first.  In  equivalent 
normal  solution  (49  grams  per  litre)  about  50  per  cent,  of  the  total 
hydrogen  is  present  in  the  ionic  condition  at  18°,  so  that  sulphuric 
acid,  though  a  very  s'trong  acid,  is  not  quite  so  strong  as  hydrochloric 
or  nitric  acids  (80  to  82  per  cent,  ionisation)  when  solutions  of  equal 
concentration  in  total  hydrogen  are  compared.  With  univalent 
metals,  sulphuric  acid  forms  salts  of  the  two  types,  M'HSO4  and 
M'2SO4;  with  bivalent  metals,  the  familiar  salts  are  of  the  type 
M"SO4  (e.g.  CuSO4  ;  BaSO^.  As  the  ion  HSO4'  gives  a  considerable 
concentration  of  H*  ions,  even  in  moderate  dilution,  the  acid  salts, 
e.g.  NaHSO4,  have  an  acid  reaction  in  solution. 

The  use  of  sulphuric  acid  in  preparing  other  acids  from  their  salts 
has  been  repeatedly  referred  to.  Its  value  in  preparing  volatile  acids 
(e.g.  hydrochloric  acid,  nitric  acid),  depends  upon  two  factors,  its 
strength  and  its  slight  volatility.  In  virtue  of  its  strength  it  displaces 


SULPHUR,    SELENIUM  AND   TELLURIUM      311 

a  considerable  proportion  of  the  other  acid  from  combination,  and 
the  mixture  can  be  raised  to  a  high  temperature  to  drive  off  the 
volatile  acid  without  much  danger  of  the  sulphuric  acid  volatilizing. 
Its  application  in  the  preparation  of  certain  other  acids  (e.g,  chloric 
acid,  p.  181),  depends  upon  an  entirely  different  property,  viz.:  the 
insolubility  of  its  calcium,  barium,  and  lead  salts  in  water.  Thus,  if 
a  calcium  salt  of  the  acid  in  question  is  treated  with  the  theoretical 
amount  of  sulphuric  acid,  the  acid  can  be  obtained  practically  pure 
by  filtering  off  the  calcium  sulphate. 

Under  certain  circumstances  sulphuric  acid  -acts  as  an  oxidizing 
agent,  being  reduced  to  sulphur  dioxide.  The  oxidation  of  carbon, 
of  sulphur,  and  of  copper  by  sulphuric  acid  have  already  been 
mentioned  in  connexion  with  sulphur  dioxide. 

The  action  of  sulphuric  acid  on  metals  depends  on  the  concentra- 
tion of  the  acid  as  well  as  on  the  nature  of  the  metal.  Dilute  acid 
readily  dissolves  iron,  zinc,  and  magnesium,  with  liberation  of  hydro- 
gen (p.  466)  : 

Zn  +  H2S  O4->ZnSO4  +  H2  f  , 

but  has  no  action  on  copper  or  lead.  Copper  is,  however,  attacked 
by  the  concentrated  acid  on  heating,  with  formation  of  copper 
sulphate  and  sulphur  dioxide  (p.  299)  : 


It  is  probable,  though  not  definitely  proved,  that  the  first  stage  in 
this  reaction  leads  to  the  formation  of  hydrogen  : 

Cu  +  H2SO4->CuSO4  +  H2, 

which  at  the  high  temperature  of  the  experiment  reduces  part  of  the 
sulphuric  acid  : 


The  action  of  concentrated  sulphuric  acid  on  mercury  and  on 
silver  is  similar  to  that  on  copper,  as  is  its  action  on  iron  : 

Fe  +  2H2SO4->FeSO4  +  SO2+2H2O. 

In  the  case  of  iron,  therefore  (and  certain  other  metals),  the  products 
differ  according  as  dilute  or  concentrated  acid  is  used. 

Sulphuric  acid  is  a  substance  of  enormous  importance  in  chemical 
industry.  It  is  used,  for  instance,  in  the  manufacture  of  sodium 
carbonate  and  of  the  more  important  mineral  acids,  in  the  manu- 
facture of  manures,  and  in  the  preparation  of  .organic  dyes. 


3i2     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Sulphates  —  The  more  important  points  in  regard  to  sulphates 
have  already  been  mentioned  in  connexion  with  the  chemical  properties 
of  sulphuric  acid.  All  sulphates  are  soluble  in  water  with  the  ex- 
ception of  those  of  calcium,  strontium,  barium  and  lead,  which  are 
practically  insoluble.  The  normal  sulphates  of  the  alkalis  are  stable 
on  heating  ;  the  acid  sulphates  lose  water  and  form  pyrosulphates 
(see  below)  : 

2NaHSO4-»Na2S2O7 


The  sulphates  of  the  heavy  metals  are  decomposed  by  heat,  the 
corresponding  oxide  and  either  sulphur  trioxide  or  (for  those  decom- 
posing at  still  higher  temperatures)  a  mixture  of  sulphur  dioxide  and 
oxygen  being  formed. 

The  insolubility  of  barium  sulphate  in  water  is  taken  advantage  of 
as  a  test  for  sulphates.  Any  solution  containing  SO4"  ions  gives,  with 
a  soluble  barium  salt,  a  white  precipitate  of  barium  sulphate,  insoluble 
in  acids. 

PYROSULPHURIC  ACID  (FUMING  SULPHURIC  ACID) 
H2S207  (H20,2S03) 

Preparation—  (i)  By  passing  sulphur  trioxide  into  concentrated 
sulphuric  acid  — 

H2SO4  +  SO3->H2S2O7. 

The  trioxide  for  this  purpose  is  made  by  the  contact  process. 

(2)  The  acid  was  formerly  prepared  by  roasting  ferrous  sulphate  in 
the  air  and  then  distilling  from  clay  retorts  (p.  560),  the  resulting 
sulphur  trioxide  being  collected  in  water  or  sulphuric  acid. 

Properties  —  The  pure  acid  is  solid  at  the  ordinary  tempera- 
ture, the  crystals  melting  at  36°,  and  it  fumes  strongly  in  the  air.  The 
commercial  "  fuming  sulphuric  acid"  consists  of  sulphuric  acid  con- 
taining varying  proportions  of  dissolved  sulphur  trioxide.  Fuming 
sulphuric  acid  is  sometimes  known  as  Nordhausen  sulphuric  acid, 
because  it  was  formerly  prepared  at  Nordhausen  in  the  Hartz 
Mountains  by  distilling  roasted  ferrous  sulphate.  Fuming  sulphuric 
acid  is  used  in  enormous  quantities  in  the  dye  industry. 

The  salts  of  pyrosulphuric  acid,  the  pyrosulphates  ,  are  best  pre- 
pared by  heating  acid  sulphates  : 


They  are  not   known  .in   aqueous   solution,  as   on   treatment  with 


SULPHUR,    SELENIUM   AND   TELLURIUM      313 

water  they  immediately  revert  to  the  acid  sulphates.      The  above 
reaction  is  therefore  reversible. 

Persulphuric  Anhydride,  S2O7,  is  obtained  in  small  drops  by  the  prolonged 
action  of  a  silent  electric  discharge  on  a  mixture  of  oxygen  and  sulphur  dioxide 
or  trioxide.  It  is  unstable,  breaking  up  slowly  even  at  room  temperature  into 
sulphur  trioxide  and  oxygen—  it  dissolves  in  water  to  form  persulphuric  acid, 


Persulphuric  Acid,  H2S2O8—  Preparation—  (i)  When  concentrated  sulphuric 
acid  (50  per  cent.)  is  cooled  and  subjected  to  electrolysis  with  a  platinum  wire 
for  anode  no  oxygen  is  evolved,  but  the  solution  round  the  anode  is  found  to 
contain  a  new  compound,  persulphuric  acid,  H2S.2O8.  Its  formation  is  readily 
understood  when  it  is  remembered  that  in  concentrated  solution  sulphuric  acid  is 
ionised  almost  exclusively  according  to  the  equation 


When  discharged  at  the  anode,  two  HSO4  groups  unite  to  form  the  acid  in 
question  : 


(2)  Persulphuric  acid  is  also  obtained  under  certain  conditions  when  a  con- 
centrated solution  of  hydrogen  peroxide  is  added  to  sulphuric  acid— 


The  reaction  is  reversible. 

(3)  The  pure  acid  can  be  obtained  in  colourless  crystals  (m.p.  60°)  by  the  action 
of  the  calculated  amount  of  100  per  cent,  hydrogen  peroxide  on  chlorosulphonic 
acid  by  the  method  described  in  connexion  with  Caro's  acid  (see  below) — 

H 

+  H202->H2S208+2HC1. 

4.  Persulphates — Persulphuric  acid  is  very  unstable  (see  below),  but  the 
corresponding  alkali  salts,  the  persulphates,  are  fairly  stable,  and  are  prepared 
by  electrolysis  of  the  acid  sulphates.  When  a  saturated  solution  of  potassium 
hydrogen  sulphate  is  subjected  to  prolonged  electrolysis  with  a  large  current, 
the  solution  being  kept  cool,  potassium  persulphate,  K2S2O8,  separates  in  colour- 
less crystals.  The  persulphuric  acid  first  formed  enters  into  double  decom- 
position with  the  excess  of  acid  sulphate  and  the  slightly  soluble  potassium 
persulphate  separates.  Other  persulphates  are  prepared  from  this  salt  or  the 
sodium  salt  by  double  decomposition. 

Properties— Persulphuric  acid  and  its  salts  in  aqueous  solution  readily  give 
up  oxygen,  forming  sulphuric  acid  and  sulphates  respectively,  and  are  therefore 
powerful  oxidizing  agents : 

2H2S2O8 + 2H2O-»4H2SO4+  O2. 

Solid  persulphates  are  moderately  stable.     Persulphates  can  be  separated  from 
sulphates  by  taking  advantage  of  the  solubility  of  barium  persulphate  in  water. 


3i4     A   TEXT-BOOK    OF    INORGANIC    CHEMISTRY 

Permonosulplniric  Acid  (Caro's  acid),  H2SO5,  is  obtained  under  certain  con- 
ditions by  the  action  of  hydrogen  peroxide  on  sulphuric  acid  : 

H2SO4  +  H202^H2S05 + H20 , 

or,  better,  by  adding  100  per  cent,  hydrogen  peroxide  to  well-cooled  chloro- 
sulphonic  acid  (p.  300).  When  the  evolution  of  hydrogen  chloride  has  ceased, 
the  reaction  mixture  is  allowed  to  warm  slowly,  and  the  remainder  of  the  hydro- 
gen chloride  removed  by  suction  : 

/OH  /OH 

S02<         +  H202-»SO/  +HC1. 

NCI  \0-OH 

Caro's  acid,  obtained  as  above,  is  a  crystalline  mass,  which  melts  at  45°.  In  aqueous 
solution  it  is  a  strong  oxidizing  agent  and  is  distinguished  from  persulphuric 
acid'by  the  fact  that  it  immediately  sets  free  iodide  from  potassium  iodide. 

Disulphur  Trioxide,  S2O3,  is  obtained  by  adding  powdered  sulphur  to  liquid 
sulphur  trioxide.  It  occurs  in  bluish-green  crystals,  readily  breaks  up  into  its 
components,  and  is  at  once  decomposed  by  water,  the  chief  products  being 
sulphuric  acid  and  sulphur. 

Hyposulphurous  Acid,  H2S2O4 — Preparation — (i)  When  zinc  acts  on  sul- 
phurous acid  in  aqueous  solution,  the  hydrogen  which  is  presumably  the  first 
product  of  the  reaction  reduces  the  sulphurous  acid  to  hyposulphurous  acid  : 

2H2S03  -f  H2->2H2O  +  H2S204, 

and  the  solution  is  found  to  have  powerful  reducing  properties. 

(2)  On  the  commercial  scale,  a  solution  containing  sodium  hyposulphite  is 
prepared  by  the  action  of  zinc  on  a  concentrated  solution  of  sodium  hydrogen 
sulphite : 

4NaHSO3  +  Zn->ZnSO3+ Na2SO3  +  Na^S^  +  2H2O, 

From  this  solution  pure  Na2S2O4  can  be  prepared  by  fractional  crystallization. 

Properties — Hyposulphurous  acid  in  aqueous  solution  is  a  yellow,  unstable 
liquid,  and,  corresponding  with  the  fact  that  it  is  the  oxyacid  of  sulphur  with 
least  oxygen,  is  a  powerful  reducing  agent.  Both  the  acid  and  its  salts  rapidly 
absorb  oxygen  from  the  air.  The  mixture  described  in  method  of  preparation 
(2)  is  used  in  the  dyeing  industry  for  reducing  indigo  to  soluble  indigo-white. 


THIOSULPHURIC  ACID,  H2S2O3 

The  acid  itself  is  extremely  unstable,  and  decomposes  immediately 
it  is  liberated  from  its  salts.  The  salts,  the  thiosulphates,  are  quite 
stable,  and  the  normal  sodium  salt,  Na2S2O3,  is  largely  used  for 
photographic  purposes  under  the  name  sodium  hyposulphite,  or 
"hypo." 

Preparation — (i)  The  thiosulphates  are  readily  obtained  by  boiling 
solutions  of  the  corresponding  sulphites  with  flowers  of  sulphur  : 

+  S->Na2S263. 


SULPHUR,    SELENIUM   AND   TELLURIUM      315 

Properties  —  Sodium  thiosulphate  forms  large  colourless  crystals, 
containing  5H2O,  which  are  readily  soluble  in  water.  When  a  free 
acid  is  added  to  a  solution  of  a  thiosulphate,  the  thiosulphuric  acid, 
which  may  be  assumed  as  the  first  product  of  the  reaction,  decom- 
poses rapidly  into  sulphur  dioxide,  water,  and  free  sulphur  : 

Na2S2O3  +  2HCl->2NaCl  +  H2S2O3 


Thiosulphates  in  aqueous  solution  act  on  the  insoluble  silver  halides 
forming  soluble  double  salts  :  AgCl  +  2Na2S2O;j-$»Na3  Ag(S2O3)2  +  NaCl. 
The  use  of  sodium  thiosulphate  in  photography  depends  upon  this  fact. 
Mild  oxidizing  agents,  such  as  iodine,  oxidize  thiosulphates  to 
tetrathionates  (qsv.}  ;  with  powerful  oxidizing  agents,  e.g.  chlorine, 
sulphates  are  formed  :  Na2S2O3  +  Cl2  +  H2O->Na2SO4 


POLYTHIONIC  ACIDS 

As  has  already  been  mentioned,  four  oxyacids  of  sulphur,  each  containing  two 
atoms  of  hydrogen  and  six  atoms  of  oxygen,  are  known.  The  names  and 
formulae  are  —  Dithionic  acid,  H2S2O6  ;  trithionic  acid,  H2S3O6;  tetrathionic 
acid,  H2S4O6  ;  and  pentathionic  acid,  H2S5O6.  The  salts  of  these  acids  are  quite 
stable  compounds.  The  respective  acids  are  usually  prepared  by  the  action  of 
sulphuric  acid  on  the  corresponding  barium  salts,  They  are  not  known  in  the 
free  state,  and  are  unstable  even  in  aqueous  solution. 

Dithionic  Acid,  H^gOg  —  The  manganese  salt  is  prepared  by  passing  sulphur 
dioxide  into  water  in  which  manganese  dioxide  is  suspended  : 

2SO2  +  MnO2->MnS2O6. 

From  this  compound  the  barium  salt,  and  then  the  fre°  acid',  can  be  obtained. 

Trithionic  Acid,  H2S3O6  —  The  potassium  salt  is  obtained  by  saturating  a 
solution  of  potassium  thiosulphate  with  sulphur  dioxide  : 

2K2S2O3  +  3SO2->2  K2S3O  +  S. 


Tetrathionic  Acid,  H^Og—  As  already  explained,  the  sodium  salt  is  obtained 
by  the  action  of  iodine  on  sodium  thiosulphate  : 

2Na2S2O3  +  I2->Na2S4O6  +  2NaI. 

Pentathionic  Acid,  II2S5OC  —  A  solution  which  contains  chiefly  pentathionic 
acid,  mixefl  with  other  polythionic  acids  in  small  proportion,  is  obtained  by 
passing  a  limited  amount  of  hydrogen  sulphide  into  a  solution  of  sulphur  dioxide  : 

5S02  +  sHjjS-^HjjSgOe  +  4H2O  +  SS. 

The  ultimate  products,  v/hen  excess  of  hydrogen  sulphide  is  used,  are  water  and 
sulphur,  as  already  stated. 


316     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

CHLORIDES  AND  AMIDES  OF  SULPHURIC  ACID 

'  Tliionyl  Chloride,  SOC12,  is  obtained  by  passing  dry  sulphur  dioxide  over 
phosphorus  pentachloride  and  separating  the  products  by  fractional  distillation  : 


It  is  a  colourless  liquid  which  boils  at  78°,  fumes  in  the  air,  and  is  immediately 
decomposed  by  water  with  formation  of  sulphurous  acid.  It  may  be  looked 
upon  as  being  derived  from  sulphur  dioxide  by  replacing  one  oxygen  by  two 
chlorine  atoms,  or  from  sulphurous  acid  by  displacement  of  two  hydroxyl  groups 
by  two  chlorine  atoms. 

Sulplmryl  Chloride,  SO2C12,  is  obtained:  (i)  by  direct  combination  of  equal 
volumes  of  sulphur  dioxide  and  chlorine  under  the  influence  of  sunlight,  or  with 
the  addition  of  camphor  as  catalytic  agent  ;  (2)  by  heating  sulphuric  acid  with 
phosphorus  pentachloride.  It  is  a  colourless  liquid,  which  boils  at  70°,  fumes 
in  contact  with  moist  air,  and  is  immediately  decomposed  by  excess  of  water  into 
a  mixture  of  sulphuric  and  hydrochloric  acids  : 

SO2C12  +  2HOH->SO2(OH)2+  sHCL 

Chlorosulphonic  Acid,  OH-SO2.C1,  is  obtained  by  the  action  of  phosphorus 
pentachloride  on  sulphuric  acid  :  SO2(O  H  )2  +  PC15->OH  •  SO2  •  Cl  +  POC13  +  HC1  , 
or  by  direct  combination  of  sulphur  trioxide  and  hydrogen  chloride.  It  is  a 
colourless  fuming  liquid  of  density  1.776,  boils  at  153°,  and  is  immediately  decom- 
posed by  water  with  formation  of  sulphuric  and  hydrochloric  acids.  It  is  evident 
that  chlorosulphonic  acid  is  intermediate  in  composition  to  sulphuryl  chloride 
and  sulphuric  acid. 

Amides  —  These  compounds  are  derived  from  sulphuric  acid  by  displacing  one 
or  both  hydroxyl  groups  by  univalent  NH2  groups.  Sulphamide  ,  SO2(NH2)2, 
and  sulphaminic  acid,  OH-SO2-NH2,  obtained  by  the  action  of  ammonia  on 
sulphuryl  chloride  and  chlorosulphonic  acid  respectively,  are  colourless  crystalline 
solids,  and  on  hydrolysis  both  give  a  mixture  of  sulphuric  and  hydrochloric 
acids. 

Graphic  Formulae  of  Sulphur  Compounds—  As  regards 
sulphuric  acid,  the  presence  of  two  hydroxyl  groups  is  shown  by  its 
relationship  to  sulphuryl  chloride  and  sulphamide,  and  the  similar 
(symmetrical)  arrangement  of  the  two  hydrogen  atoms  is  further  sup- 
ported by  the  fact  that  the  same  compound  is  obtained  when  one  or 
other  of  them  is  displaced  by  an  organic  group  such  as  C2H6.  The 
formula  of  the  acid  is  therefore  SO2(OH)2,  and  might  be  either 
CK,  /OH  O\  /OH 

\S\  or    I    J>S<^         ,  according  as  sulphur  is  assumed  to  be 

O^    \OH         O/       \OH 
sexavalent   or   quadrivalent.     The   former  alternative   is   the  more 


SULPHUR,  SELENIUM    AND  TELLURIUM      317 

probable,  and  receives  support  from  the  existence  of  sulphur  hexa- 

fluoride. 

/OH 

As  regards  sulphurous  acid,  two  formulae  are  possible,  O  =  S\ 

\OH 

O^      /H 

and       ;Z5\         ,  and  as  organic  derivatives  corresponding  with  each 
O<^     \OH 

/  >OC2H6  O^     /C2H5\ 

are  known  I  e.g.  Q  —  S\  and       \SC  the    acid   may 

\  \OH  O^   \OH  / 

perhaps  be  an  equilibrium  mixture  of  both  forms.  The  relationship 
to  thionyl  chloride  lends  support  to  the  first  formula;  the  slight 
ionisation  of  the  second  hydrogen  atom  (compare  hydrogen  sulphide) 
is  perhaps  more  in  accord  with  the  unsymmetrical  formula. 

Crystallography — In  the  foregoing  chapters  it  has  been  fre- 
quently mentioned  that  many  substances  can  be  obtained  in  definite 
geometrical  forms,  so-called  crystals.  Crystals  may  be  obtained  on 
allowing  a  solution  of  the  substance  to  cool  (e.g.  sulphur  in  carbon 
disulphide),  by  allowing  a  fused  substance  to  solidify  (e.g.  crystals  of 
prismatic  sulphur),  by  allowing  the  vapour  of  a  substance  to  condense 
(e.g.  iodine),  and  in  other  ways.  When  a  substance  occurs  in  a  non- 
crystalline  form  it  is  said  to  be  amorphous.  Sometimes  the  same 
substance  may  exist  in  a  crystalline  and  in  an  amorphous  form  (e.g. 
sulphur) ;  on  the  other  hand,  some  substances  occur  invariably  in  the 
amorphous  form. 

The  most  marked  feature  of  a  crystal  is  its  regularity  of  form.  The 
great  majority  of  well-developed  crystals  show  some  kind  of  symmetry. 
From  this  point  of  view  the  possible  type  of  crystals  are  considered 
with  reference  to  imaginary  points,  planes,  and  axes,  termed  respec- 
tively points  of  symmetry,  planes  of  symmetry,  and  axes  or  lines  of 
symmetry. 

A  crystal  has  a  point  or  centre  of  symmetry  when  for  each  face  on 
the  crystal  there  is  another  parallel  to  it  on  the  other  side  of  the 
crystal. 

A  crystal  has  a  plane  of  symmetry  when  an  imaginary  plane  can  be 
drawn  through  it  so  that  the  faces  are  in  pairs  symmetrically  placed 
with  reference  to  the  plane ;  in  other  words,  the  plane  cuts  the  crystal 
into  two  parts,  one  of  which  is  the  mirror  image  of  the  other.  Crystals 
may  have  from  o  to  9  planes  of  symmetry. 

A  crystal  has  an  axis  of  symmetry  when  it  can  be  rotated  about  an 
imaginary  line  (drawn  through  the  centre  of  the  crystal)  in  such  a 


3i8     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


way  that  each  face  occupies  the  position  previously  taken  by  another, 
at  least  once  in  a  rotation  through  360°.  During  the  rotation 
through  360°  this  periodic  recurrence  of  the  same  aspect  may  take 
place  either  two,  three,  four,  or  six  times,  and  the  axes  are  termed 
binary,  trigonal,  tetragonal,  and  hexagonal  respectively. 

All  known  crystals  may  be  classified  in  six  systems,  based  upon  the 
relations  of  their  crystallographic  axes. 

(1)  The  cubic  (regular  or  isometric)  system — Crystals  belonging  to 
this  system  can  be  referred  to  three  axes  of  equal  length,  all  at  right 
angles  to  each  other.     The  regular  cube  (examples  :  common  salt, 
fluor-spar  (Fig.  60)),  and  the  octahedron  (example,  the  alums  (Fig.  61)) 
belong  to  this  class. 

(2)  The  hexagonal  system — Crystals   belonging   to  this  class  are 
referred  to  four  axes,  three  of  which,  in  one  plane,  are  equal  in  length 


FIG.  60. 


FIG.  61. 


FIG.  62. 


FIG.  63. 


and  intersect  at  angles  of  120° ;  the  fourth,  known  as  the  principal 
axis,  is  at  right  angles  to  the  plane  of  the  other  three,  and  differs 
from  them  in  length.  Ice,  quartz,  calc-spar  (Fig.  62),  and  many  other 
important  substances  belong  to  this  class. 

(3)  The  tetragonal  system — The  crystals  are  referred  to  three  axes, 
of  which  two  only  are  of  equal   length.     Example :  stannic  oxide 
(Fig.  63). 

(4)  The  orthorhombic  (rhombic)  system — The  crystals  are  referred 
to  three  axes  at  right  angles,  unequal  in  length.     Examples  :  rhombic 
sulphur  (Fig.  64),  potassium  nitrate,  magnesium  sulphate. 

(5)  The  monoclinic  (monosymmetric)  system — The  crystals  belong- 
ing to  this  system  are  referred  to  three  axes  of  unequal  length  ;  two 
of  these  are  not  at  right  angles  ;  the  third  is  at  right  angles  to  the 
plane  of  the  other  two.     Examples :  monoclinic  sulphur  (Fig.  65), 
ferrous  sulphate  heptahydrate,  gypsum  (Fig.  66). 

(6)  The  triclinic  (asymmetric)  system — The  crystals  belonging  to 


SULPHUR,    SELENIUM   AND   TELLURIUM      319 

this  system  are  referred  to  three  axes  of  unequal  length,  intersecting 
one  another  at  a  point  at  angles  which  are  unequal  and  not  right  angles. 
There  are  two  divisions  of  this  system  ;  in  one  the  crystals  have  a 
centre  of  symmetry  only,  in  the  other  division  there  is  no  symmetry. 
Examples  :  copper  sulphate  pentahydrate  (Fig.  67),  potassium  bi- 
chromate. 

The  more  important  terms  relating  to  crystals  have  already  been 


FIG.  64.  FIG.  65.  FIG.  66.  FIG.  67. 

mentioned.  Where  the  same  substance  occurs  in  different  crystalline 
forms  it  is  said  to  be  polymorphous ;  if  only  in  two  forms,  dimorphous. 
Substances  of  the  same  crystalline  form  are  said  to  be  isomorphous 
(p.  120).  It  sometimes  happens  that  a  dimorphous  substance  is  iso- 
morphous with  another  dimorphous  substance  in  both  its  forms  ;  this 
is  termed  isodinwrphism. 

SELENIUM  AND  TELLURIUM 

In  the  group  of  which  oxygen  and  sulphur  are  the  first  elements  there  are  two 
other  elements,  selenium  (atomic  weight— 79.1)  and  tellurium  (atomic  weight= 
127.6),  which  show  many  analogies  with  sulphur,  and  are  most  conveniently 
considered  here.  Both  form  compounds  with  hydrogen,  HgSe  and  H2Te,  of  the 
same  type  as  hydrogen  sulphide.  Selenium  gives  an  oxide,  SeO2,  and  tellurium 
two  oxides,  TeO2  and  TeO3,  of  the  same  type  as  the  familiar  oxides  of  sulphur. 
It  will  be  shown,  however,  that,  unlike  the  oxides  of  sulphur,  those  of  tellurium 
have  basic  as  well  as  acidic  properties.  Two  oxyacids,  HoSeOs,  tellurous  acid, 
and  H2TeO4,  telluric  acid,  are  known. 

SELENIUM 
Symbol,  Se.     Atomic  weight =79.1.     Molecular  weight =158. 2. 

Occurrence — Selenium,  though  a  fairly  widely-distributed  element,  is  met  with 
only  in  small  quantities.  It  occurs  in  combination  with  certain  metals,  especially 
lead,  mercury ,  and  copper,  in  the  Hartz  Mountains ;  but  the  chief  sources  are  certain 
pyrites  (FeS2)  in  which  the  sulphur  is  partly  displaced  by  selenium.  When  these 
pyrites  are  used  in  the  manufacture  of  sulphuric  acid  the  selenium  is  converted 
into  the  dioxide,  and  is  partly  deposited  in  the  flues  of  the  pyrite-burners  and 
partly  carried  forward  into  the  chambers,  where  it  is  reduced  to  selenium  by 


320     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

sulphur  dioxide,  and  collects  in  the  mud  on  the  floor  of  the  chamber.  In  this 
"  chamber  mud  "  selenium  was  discovered  in  1817  by  Berzelius.  The  name  (from 
<re\~f)vy,  the  moon)  was  given  on  account  of  its  analogy  with  tellurium  (from 
tellus,  the  earth),  discovered  not  long  before. 

Preparation  —  Selenium  is  obtained  from  the  flue  dust  or  from  the  chamber  mud 
by  heating  first  with  nitric  acid,  in  order  to  convert  it  completely  into  seleriic  acid. 
The  latter  is  then  boiled  with  hydrochloric  acid,  whereby  it  is  reduced  toselenious 
acid,  H2SeO3,  with  liberation  of  chlorine  ;  the  selenious  acid  is  then  reduced  by 
means  of  sulphur  dioxide  to  selenium,  which  separates  in  red  flakes: 

H2Se03  +  2SO2  +  H20-»Se  +  2H2SO4. 

Properties—  Selenium  exists  in  at  least  three  allotropic  modifications:  (i)An 
amorphous  red  form,  obtained,  for  example,  by  reducing  selenious  acid  with 
sulphur  dioxide.  It  is  soluble  in  carbon  disulphide,  and  separates  from  solution 
as  a  second  modification  —  (2)  Red  crystalline  selenium,  melting  at  170° 
to  180°.  (3)  A  third  modification,  termed  metallic  selenium,  is  obtained  by 
heating  amorphous  selenium  to  97°,  or  by  quickly  cooling  melted  selenium 
to  210°  and  keeping  for  some  time  at  that  temperature.  It  occurs  in  grey 
crystals,  is  insoluble  in  carbon  disulphide,  and  conducts  electricity.  It  is  a 
remarkable  fact  that  the  electrical  conductivity  of  metallic  selenium  is  greatly 
increased  on  exposure  to  light.  According  to  recent  investigations,  metallic 
selenium  consists  of  two  forms  of  very  different  conducting  power  in  equilibrium, 
and  the  equilibrium  is  displaced  by  light. 

Selenium  melts  at  217°  and  boils  at  680°.  Above  1400°  its  vapour  density  corre- 
sponds with  the  formula  Se.,  ;  at  lower  temperatures  the  molecule  is  more 
complex. 

Hydrogen  Selenide,  HaSe—  Preparation—  (i)  By  passing  hydrogen  over  sele- 
nium heated  to  400°  : 


(2)  By  acting  on  ferrous  selenide  with  hydrochloric  acid  : 
FeSe  +  2HCl-»FeCl2  +  HgSe. 

Properties  —  Hydrogen  selenide  is  a  colourless  gas  with  an  odour  like  horse- 
radish, and  is  more  poisonous  than  hydrogen  sulphide.  It  is  moderately  soluble 
in  water,  and  the  aqueous  solution  deposits  selenium  on  exposure  to  the  air  owing 
to  oxidation.  When  passed  through  solutions  of  salts  of  the  heavy  metals  the 
selenides,  being  insoluble,  are  precipitated. 

Selenium  Monochloride,  SeaCLj,  and  selenium  tetrachloride,  SeCl4,  are 
formed  by  direct  combination  of  the  elements.  The  former  is  a  brownish-yellow 
oily  liquid  ;  the  latter  a  white  crystalline  solid  which  can  be  sublimed  without 
decomposition,  but  undergoes  partial  dissociation  into  its  elements  when  heated 
above  200°. 

OXIDES  AND  OXYACIDS  OF  SELENIUM 

Selenium  Dioxide,  SeO2,  the  only  known  oxide  of  selenium,  is  obtained  by 
burning  selenium  in  the  air.  It  occurs  in  long  white  needles,  which  sublime 
(without  previously  melting)  at  310°.  It  dissolves  in  water,  forming  selenious 
acid,  H2SeO,. 


SULPHUR,   SELENIUM   AND   TELLURIUM      321 

Selenious  acid,  H2SeO3,  is  formed  when  selenium  dioxide  is  dissolved  in  water, 
and  unlike  sulphurous  acid  can  be  isolated  by  evaporating  the  solution,  when  it 
separates  in  colourless  prismatic  crystals.  It  is  a  dibasic  acid,  forming  normal 
and  acid  salts.  Reducing  agents,  for  example,  sulphur  dioxide  or  stannous 
chloride,  reduce  it  to  selenium  : 


Selenlc  Acid,  HgSeOj,  >s  obtained  by  the  action  of  chlorine  (or  bromine)  on 
selenious  acid  : 

HsSeOg  +  C12  +  H2CqtH2Se04  +  2HC1. 

The  action  is  reversible,  as  selenic  acid  can  be  reduced  to  selenious  acid  by 
boiling  with  hydrochloric  acid.  Selenic  acid  is,  therefore,  a  more  powerful 
oxidizing  agent  than  sulphuric  acid,  which  oxidizes  hydrobromic  acid  (p.  156),  but 
not  hydrochloric  acid.  The  pure  acid  forms  colourless  crystals,  which  melt  at 
58°.  The  95  per  cent,  solution  is  an  oily  liquid,  similar  to  sulphuric  acid;  its 
density  is  2.6.  Barium  selenate  is  as  insoluble  in  water  as  barium  sulphate. 

TELLURIUM 
Symbol,  Te.     Atomic  Weight,  127.6 

Occurrence  —  Tellurium  is  a  rare  element.  It  occurs  naturally  to  a  small 
extent  in  the  free  condition,  but  more  commonly  in  combination  with  silver,  gold, 
bismuth,  and  other  metals  in  Transylvania,  Hungary,  California,  Brazil,  and 
Bolivia. 

The  preparation  of  tellurium  from  its  ores  is  a  complicated  process. 

Properties—  Tellurium,  like  the  other  elements  of  this  group,  occurs  in 
different  allotropic  modifications.  When  precipitated  from  solution,  a  black, 
amorphous  form  is  obtained,  but  on  heating  it  fuses  and  solidifies  as  a  silvery- 
white  substance  with  metallic  lustre;  its  density  is  6.24  and  it  melts  at  452°. 
Metallic  tellurium  conducts  heat  and  electricity,  but  is  not  a  good  conductor. 

Hydrogen  telluride,  H2Te,  is  obtained  by  the  action  of  hydrochloric  acid  on 
zinc  telluride  : 

2Tef. 


Properties  —  Hydrogen  telluride  is  a  colourless  gas  with  a  disagreeable  odour, 
and  is  very  poisonous.  It  is  fairly  soluble  in  water,  and  the  solution  deposits 
tellurium  on  exposure  to  air.  When  passed  into  solutions  of  salts  of  the  heavy 
metals,  the  corresponding  tellurides  are  precipitated.  Hydrogen  telluride  is 
more  easily  decomposed  into  its  elements  by  heat  than  the  corresponding  sulphur 
and  selenium  compounds. 

Tellurium  dichloride,  TeCl2,  and  Tellurium  tetrachloride,  TeCl4,  are  ob- 
tained by  direct  combination  of  the  elements.  The  former  occurs  in  small,  nearly 
black  crystals,  the  latter  forms  colourless  crystals.  Both  can  be  volatilized  with- 
out decomposition  at  high  temperatures.  They  are  decomposed  by  water,  the 
latter  according  to  the  equation 

TeCl4+3H20^>H2Te03+4HCl, 

but  the  reaction  is  reversed  when  the  hydrochloric  acid  concentration  is  con- 
siderable. 

21 


322     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


OXIDES  AND  OXYACIDS  OF  TELLURIUM 

Tellurium  dioxide,  TeO2,  is  obtained  by  burning  tellurium  in  the  air  or  by 
heating  tellurous  acid.  It  is  a  white  crystalline  powder,  slightly  soluble  in  water, 
and  is  at  the  same  time  an  acidic  and  a  basic  oxide  (see  below). 

Tellurous  acid,  H2TeO3,  is  obtained  by  oxidizing  tellurium  with  nitric  acid. 
It  occurs  as  a  white  powder,  slightly  soluble  in  water.  It  is  a  weak  dibasic  acid, 
forming  normal  and  acid  salts  with  the  alkalis  which  are  soluble  in  water.  On 
the  other  hand,  it  forms  salts  with  strong  acids,  which  may  be  looked  upon  as 
being  derived  from  the  diacidic  base,  TeO(OH)2,  or  from  the  tetracidic  base, 
Te(OH)4—  that  is,  H2TeO3,H2O.  Both  types  of  salt  are  considerably  hydrolyzed 
in  solution,  in  accordance  with  the  rule  that  when  the  same  substance  has  both 
basic  and  acidic  properties  it  is  invariably  weak,  both  as  base  and  as  acid. 

Tellurium  trioxide,  TeO8,  is  obtained  by  heating  telluric  acid  above  160°  : 

H2TeO4->H2O  +  TeO3. 

It  is  a  yellow  powder,  practically  insoluble  in  water,  and  its  acidic  properties  are 
extremely  weak.     It  splits  up  into  the  dioxide  and  oxygen  on  heating  strongly. 

Telluric  acid,  H2TeO4,  is  obtained  by  the  action  of  powerful  oxidizing  agents, 
for  example,  chromic  acid,  on  tellurous  acid.  It  is  most  conveniently  obtained 
by  fusing  tellurium  or  tellurium  dioxide  with  potassium  carbonate  and  nitrate  : 


Barium  tellurate  is  then  obtained  from  the  potassium  salt  by  double  decomposi- 
tion with  barium  chloride  and  treated  with  the  calculated  amount  of  dilute 
sulphuric  acid  : 

BaTeO4  +  H2SO4->BaSO4  +  H2TeO4. 

On  filtering  and  evaporating  the  solution,  telluric  acid  separates  in  the  form  of 
crystals  of  the  composition  H2TeO4,2H2O. 

Telluric  acid  dihydrate,  H2TeO4,2H2O,  or  Te(OH)6,  is  a  crystalline  powder, 
very  slightly  soluble  in  water;  its  acidic  properties  are  extremely  weak,  and  it 
therefore  .shows  practically  no  analogy  to  sulphuric  acid.  Like  tellurous  acid,  it 
has  weak  basic  properties. 

Summary  of  the  Oxygen  Group  —  Sulphur,  selenium,  and 
tellurium,  along  with  oxygen,  constitute  a  family  or  group  of  elements 
which  are  very  similar  in  chemical  behaviour,  and,  like  the  halogens, 
show  a  gradual  variation  in  physical  and  chemical  properties  with 
increase  in  atomic  weight.  The  latter  statement  is  illustrated  as 
regards  a  few  physical  properties,  in  the  accompanying  table  (cf. 
halogens,  p.  190):  — 

As  in  the  case  of  the  halogens,  however,  there  is  a  much  closer 
resemblance  between  the  last  three  elements  among  themselves  than 
between  these  elements  and  oxygen.  Sulphur  and  oxygen  resemble 
each  other  in  being  bivalent,  so  that  their  compounds  with  hydrogen 


SULPHUR,   SELENIUM   AND   TELLURIUM     323 

and  with  the  metals  are  of  the  same  type,  and  are,  in  many  cases, 
similar  in  behaviour. 


Property. 

Oxygen. 

Sulphur. 

Selenium. 

Tellurium. 

Atomic  Weight  . 

16 

32 

79.2 

127.5 

Melting-point 

-227° 

114.5° 

217° 

452° 

Boiling-point 

-181.5° 

445° 

680° 

1400° 

Density 

i.i3(atb.-pt.) 

1.96-2.05 

4.8 

6.24 

The  last  three  elements  are  divalent,  quadrivalent  and  sexavalent 
in  their  compounds,  and  therefore  the  compounds  are  of  similar  type, 
as  shown,  for  example,  in  the  hydrides  H2S,  H2Se  and  H2Te  ;  the 
oxides  SO2  and  SO3,  SeO2,  TeO2,  and  TeO3,  and  the  acids  of  the  two 
types  H2EO3  and  H2EO4  (E  =  element).  What  is  much  more  striking, 
however,  is  that  the  similarity  extends  to  the  chemical  behaviour  of 
the  compounds,  as  shown  in  detail  in  the  foregoing  pages.  The 
affinity  for  hydrogen  diminishes  with  increase  of  atomic  weight,  as 
shown  by  the  relative  stability  of  the  hydrides.  The  affinity  for 
oxygen  also  diminishes,  but  the  affinity  for  chlorine  increases,  with 
increasing  atomic  weight.  Perhaps  the  most  important  point,  how- 
ever, is  that  the  acidic  character  of  the  oxides  diminishes  with 
increasing  atomic  weight,  and  the  oxides  of  tellurium  show  also  weak 
basic  properties.  In  this  respect,  as  also  in  its  appearance  and  in  the 
relative  stability  of  its  chloride  towards  water,  tellurium  approaches 
the  metals. 


CHAPTER   XXII 
CARBON 

CARBON 

Symbol,  C.          Atomic  weight=i2.  Molecular  weight  unknown. 

Occurrence — Carbon  occurs  free  in  nature  in  two  crystalline 
modifications,  diamond  and  graphite,  and  also  in  the  amorphous 
form  as  charcoal.  In  combination  with  hydrogen,  it  occurs  in  marsh 
gas  and  in  petroleum.  In  combination  with  oxygen,  it  occurs  as 
carbon  dioxide,  which  is  an  invariable  constituent  of  the  atmosphere 
(p.  205),  and  is  found  in  maiiy  natural  waters  (p.  59) ;  and  also  as  car- 
bonates. Calcium  carbonate  is  met  with  as  chalk,  marble  and  limestone 
in  enormous  quantities.  The  carbonates  ot  magnesium,  zinc,  barium 
and  of  other  elements  also  occur  naturally. 

Carbon  is  an  essential  constituent  of  plants  and  animals.  Coal  is 
formed  as  the  result  of  the  slow  decay  of  vegetable  matter,  and 
consists  mainly  of  free  carbon,  but  other  substances  are  always  present 
in  greater  or  less  amount  (p.  330). 

ALLOTROPIC  MODIFICATIONS 

As  already  indicated,  carbon  occurs  in  two  crystalline  forms, 
diamond  and  graphite,  and  also  as  amorphous  carbon.  Under  the 
latter  name  are  included  all  the  numerous  forms  of  carbon  which  have 
no  definite  crystalline  form.  The  proof  that  all  these  substances  are 
composed  of  the  same  element  is  that  carbon  dioxide  is  the  sole  pro- 
duct of  their  combustion  in  air. 

DIAMOND 

Occurrence — Diamonds  were  originally  obtained  solely  from 
India,  being  found  in  alluvial  deposits.  They  have  since  been 
discovered  in  Brazil  (about  1727),  Australia,  South  Africa  (1867), 
and  also  in  the  United  States.  They  have  also  been  found  in 
meteorites. 

3*4 


CARBON  325 

Diamonds  vary  much  in  size,  and  are  usually  comparatively  small. 
A  remarkable  exception  is  the  Cullinan  diamond,  discovered  in  the 
Premier  mine  at  Kimberley  in  1905 ;  it  originally  weighed  over 
600  grams. 

Preparation  of  Artificial  Diamonds — After  it  was  recog- 
nized that  diamonds  are  simply  crystallized  carbon,  many  attempts 
were  made  to  prepare  them  artificially.  Great  difficulties  were, 
however,  met  with,  but  the  problem  was  to  some  extent  solved  by 
Moissan.  When  carbon  is  dissolved  in  melted  iron,  and  the  fused 
mass  allowed  to  cool,  the  carbon  separates  in  the  form  of  graphite. 
Moissan  modified  the  method  by  heating  the  mixture  of  carbon  and 
iron  to  3000°  in  the  electric  furnace  (see  below),  and  then  suddenly 
immersing  in  cold  water  the  carbon  crucible  containing  the  fluid  mass. 
A  solid  crust  forms  first  on  the  surface,  and  as  iron  saturated  with 
carbon  expands  on  solidification,  an  enormous  pressure  is  thus  exerted 
on  the  interior  partially  liquid  portion.  The  iron  was  finally  dis- 
solved away  by  acid,  and  in  the  residue,  which  consisted  mainly  of 
graphite,  minute  crystalline  particles,  having  all  the  properties  of 
diamonds,  were  found.  Some  of  the  diamonds  were  transparent, 
others  were  black.  The  largest  were  only  about  0.5  mm.  in  diameter.1 

The  mode  of  formation  of  diamonds  in  nature  is  not  thoroughly 
understood,  but  they  are  probably  formed  by  crystallization  of  carbon 
in  the  interior  of  the  earth  at  high  temperatures  under  enormous 
pressure. 

Properties — The  diamond  occurs  in  crystals  belonging  to 
the  regular  system,  the  octahedral  and  cubic  forms  being  most 
common.  When  pure  it  is  colourless  and  transparent,  but  is  very 
frequently  coloured  by  traces  of  impurities.  When  the  colour  is 
agreeable,  such  as  blue,  red  or  green,  the  stones  may  be  ^s  valuable 
for  gems  as  the  colourless  variety.  The  black  variety,  known  as  car- 
bonado or  bort,  contains  up  to  2  per  cent,  of  impurities,  and  is  useless 
as  a  gem  ;  but  on  account  of  its  extreme  hardness  is  used  for  boring 
rocks,  for  cutting  glass,  and,  in  the  form  of  a  powder,  for  cutting  and 
polishing  precious  stones,  including  diamonds.  Diamond  is  the 
hardest  substance  known.  It  has  a  very  high  refractive  index,  and  it 
is  this  property  of  scattering  light  possessed  by  the  diamond  to  such 
a  pre-eminent  extent  which  renders  it  so  valuable  as  a  jewel. 

The  average  density  is  about  3.5,  but  that  of  the  black  variety  is 
distinctly  smaller,  from  3.0  to  3.4. 

1  There  is  reason  to  suppose  that  Moissan's  artificial  diamonds  were  not  entirely 
free  from  silicon. 


326    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

The  diamond  does  not  conduct  electricity.  It  is  not  attacked  by 
acids.  When  heated  in  absence  of  air  to  temperatures  above  1000°, 
it  changes  to  graphite.  When  heated  in  air  or  oxygen,  it  readily  burns 
to  carbon  dioxide,  as  was  first  shown  by  Lavoisier  (1772). 


GRAPHITE  (BLACKLEAD) 

Occurrence — This  allotropic  modification  of  carbon  is  widely 
distributed  in  nature,  being  found  in  Bohemia,  Ceylon,  Spain,  Siberia, 
California  and  Borrowdale  in  Cumberland. 

Modes  Of  Formation — (i)  As  already  stated,  graphite  is  ob- 
tained when  diamond  is  heated  to  1000°  in  absence  of  air.  Further, 
when  carbon  is  dissolved  in  melted  iron,  it  separates  in  the  form  of 
graphite  on  cooling. 

(2)  As  natural  graphite  is  generally  very  impure,  and  the  substance 
is  of  considerable  commercial  importance,  the  discovery  by  Acheson 
of  a  cheap  method  of  obtaining  it  from  charcoal  (coal  or  coke)  proved 
of  great  value.  The  coke  or  finely  divided  coal  is  heated  with  a 
mixture  of  oxides  (including  those  of  iron  and  calcium)  in  the  electric 
furnace.  Carbides  are  first  formed  ;  these  at  the  high  temperature 
decompose  into  graphite,  which  is  deposited  in  a  very  pure  condition, 
and  the  respective  metals,  which  distil  over  into  the  colder  parts  of  the 
furnace. 

Properties — Graphite  is  a  soft,  shiny,  grayish-black  substance, 
with  metallic  lustre.  The  crystals  are  six-sided  leaflets  belonging 
to  the  monoclinic  system.  Its  density  varies  from  2.17  to  2.32. 
Graphite,  unlike  diamond,  is  a  good  conductor  of  heat  and  electricity. 
Most  natural  graphite  is  very  impure,  containing  from  40  to  70  per  cent, 
of  carbon,  and  leaving  a  large  proportion  of  ash  when  burned. 

Graphite  has  the  property  of  breaking  off  smooth,  thin  scales  on 
rubbing  or  pressing,  and  on  this  is  based  its  use  for  making  lead 
pencils,  and  for  lubricating  purposes.  In  making  pencils,  the  finely 
powdered  natural  or  artificial  graphite  is  mixed  with  a  little  clay  as 
binding  material,  and  the  semisolid  mass  formed  into  narrow  threads 
by  squeezing  it  through  a  small  opening.  It  was  formerly  supposed 
to  contain  lead,  hence  the  common  names  plumbago  and  blacklead 
often  applied  to  it. 

Graphite  is  not  readily  attacked  by  air  or  even  by  oxygen  at  high 
temperatures,  and  is  not  affected  by  the  great  majority  of  chemical 
compounds.  For  this  reason,  it  is  used,  along  with  fireclay,  in  making 
the  so-called  plumbago  crucibles.  Further,  on  account  of  its  refractory 


CARBON  327 

nature  and   its   high    conducting   power,   it   is   largely  used  as   an 
electrode  material  for  batteries  and  in  electrotyping. 

When  finely  divided  graphite  is  treated  with  potassium  chlorate  and 
either  nitric  acid  or  a  mixture  of  nitric  and  sulphuric  acids  a  gray, 
apparently  crystalline  substance,  the  so-called  "graphitic  acid,"  is 
formed.  The  nature  of  "  graphitic  acid  "  is  not  understood,  and  it  is 
doubtless  a  complicated  mixture.  On  heating,  it  decomposes  with 
explosion,  and  a  black  substance,  "  pyrographitic  acid,"  remains 
behind. 

AMORPHOUS  CARBON 

This  term  includes  all  the  varieties  of  carbon  which  are  non- 
crystalline.  The  more  important  are — charcoal,  including  among 
other  kinds  wood  charcoal  and  animal  charcoal,  also  lamp-black  or 
soot,  coke  and  gas  or  retort  carbon.  The  majority  of  these  sub- 
stances are  prepared  by  heating  carbon  compounds  in  the  absence  of 
air.  They  are  all  more  or  less  impure.  The  proportion  of  carbon 
varies  from  about  10  per  cent,  in  animal  charcoal  to  nearly  100  per 
cent,  in  carefully  purified  sugar  charcoal. 

Charcoal  is  the  general  name  applied  to  amorphous  carbon 
obtained  by  heating  substances  (other  than  coal)  containing  carbon 
in  the  absence  of  air,  that  is,  by  destructive  distillation  (p.  214).  An 
important  variety  is  wood  charcoal.  Wood  consists  essentially  of 
chemical  compounds  containing  carbon,  hydrogen  and  oxygen,  and 
the  process  of  carbonization  consists  in  the  more  or  less  complete 
removal  of  the  hydrogen  and  oxygen,  along  with  part  of  the  carbon, 
as  water,  compounds  of  carbon  and  hydrogen  and  a  great  variety  of 
other  products.  This  may  be  effected  by  destructive  distillation,  and 
also  takes  place  when  vegetable  matter  is  allowed  to  decay  in  absence 
of  air  (see  coal,  p.  330).  The  method  of  carbonizing  wood  differs 
according  as  the  other  materials,  including  acetic  acid,  tar,  etc., 
formed  in  the  process  are  required  or  not.  In  the  former  case,  the 
wood  is  placed  in  retorts,  which  are  heated  externally,  and  no  air  is 
admitted.  The  volatile  products  which  are  gases  at  the  ordinary 
temperature  are  allowed  to  escape,  and  the  others  are  condensed. 
Wood  charcoal  remains  behind  in  the  retort.  The  more  wasteful 
process  in  which  the  products  are  allowed  to  escape  is  carried  on  in 
the  forests  by  charcoal-burners.  The  wood  is  piled  in  heaps,  covered 
with  sods  and  earth,  and  ignited.  The  entry  of  air  is  carefully 
regulated  so  that  the  wood  smoulders  away  and  is  finally  completely 
carbonized.  This  process  is  also  mainly  one  of  destructive  distillation, 


328    A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the   requisite   heat  being   obtained   by   combustion   of  part   of  the 
material. 

Wood  charcoal  retains  the  shape  of  the  substance  from  which  it  is 
formed.  The  chief  impurities  are  the  mineral  substances  always 
present  in  wood,  and  also  gases.  Wood  charcoal  has  a  high  absorp- 
tive power  for  gases,  a  property  shared  by  some  of  the  other  forms  of 
carbon. 

A  much  purer  charcoal  is  obtained  by  carbonizing  substances  which 
contain  no  mineral  matter,  such  as  cane  sugar.  The  charcoal  ob- 
tained from  the  latter  substance,  after  igniting  in  a  stream  of  chlorine 
to  remove  absorbed  hydrogen,  is  practically  pure  amorphous  carbon. 

Animal  Charcoal  is  obtained  by  the  carbonization  of  bones 
in  closed  iron  retorts.  It  contains  only  about  8  to  12  per  cent,  of 
carbon,  up  to  80  per  cent,  of  calcium  phosphate,  and  some  calcium 
carbonate  and  sulphate,  but  owing  to  its  very  porous  character  has 
great  absorbing  power  for  gases,  colouring  matters,  etc.  By  treating 
it  with  acids,  nearly  pure  charcoal  can  be  obtained. 

Lamp-black  or  Soot  is  obtained  by  burning  substances  rich 
in  carbon,  such  as  petroleum,  turpentine  and  naphthalene  in  a  limited 
supply  of  air.  When  freed  from  hydrogen  and  other  compounds  by 
heating  in  a  stream  of  chlorine,  it  is  a  very  pure  form  of  carbon. 
It  is  used  in  making  paints  and  Indian  ink,  etc. 

Coke — When  coal  is  subjected  to  destructive  distillation,  gases, 
tar  and  other  products  volatilize,  and  coke  is  left  behind  in  the 
retort.  Coke  is  very  impure,  containing  most  of  the  mineral  matter 
which  forms  the  ash  of  coal,  as  well  as  sulphur  ;  the  proportion 
of  carbon  in  coke  may  reach  80  per  cent.  It  is  largely  used  for 
metallurgical  purposes. 

Gas  or  Retort  Carbon  is  found  as  a  hard  deposit  lining 
the  walls  of  retorts  used  in  the  destructive  distillation  of  coal.  It 
results  from  the  decomposition  of  volatile  products  of  distillation  on 
the  hot  walls  of  the  retorts  and  is  a  fairly  pure  product,  containing 
only  i  to  3  per  cent,  of  ash.  Its  density  is  about  2.0,  it  is  a  good 
conductor  of  electricity,  and  is  largely  used  for  making  carbon 
electrodes. 

Physical  Properties  of  Charcoal.  Adsorption- 
Carbon  may  be  regarded  as  infusible.  At  the  temperature  of  the 
electric  arc,  3500°,  it  is  partly  volatilized,  and  condenses  in  the  form 
of  graphite. 

The  properties  of  charcoal  depend  greatly  upon  its  method  of 
preparation  and  on  its  degree  of  purity.  For  example,  the  density 


CARBON 


32Q 


01  purified  lamp-black  under  ordinary  conditions  is  about  1.78,  but 
can  be  raised  to  1.87  by  prolonged  heating  in  absence  of  air.  The. 
fact  that  wood-charcoal  floats  in  water  is  due  to  the  presence  of 
absorbed  air ;  when  the  latter  is  pumped  out  the  charcoal  sinks. 
The  average  density  of  wood  charcoal  is  about  1.5  ;  of  sugar 
charcoal  about  1.8.  It  has  already  been  pointed  out  that  whereas 
gas  carbon  and  graphite  are  good  conductors  of  electricity,  diamond 
and  ordinary  charcoal  are  practically  non-conductors.  It  is  pro- 
bable that  the  conductivity  depends  on  the  presence  of  impurities. 

A  remarkable  property  of  charcoal  is  its  power  of  taking  up  gases 
and  vapours,  and  also,  to  a  less  extent, 
colouring  matters  and  other  substances 
from  solution.  This  phenomenon  is  now 
called  adsorption.  It  is  a  surface  effect 
(a  physical,  not  a  chemical  action),  and  is 
very  pronounced  for  charcoal  owing  to  the 
very  great  extent  of  surface  for  a  com- 
parative small  mass.  It  is  very  well 
shown  by  heating  a  piece  of  wood  char- 
coal to  drive  out  the  condensed  air,  and 
then  passing  it  up  into  ammonia  gas 
confined  in  a  tube  over  mercury  (Fig.  68) ; 
the  mercury  will  be  observed  to  rise  fairly 
rapidly  in  the  tube. 

The  amount  of  adsorption  is  greatest 
for  the  most  readily  liquefiable  gases,  it 
increases  with  increase  of  pressure, 
although  Henry's  law  is  not  followed,  and 
is  increased  by  lowering  the  temperature. 
Titoff  found  that  I  gram  of  cocoa-nut  charcoal  adsorbed,  at  o°  and 
76  cm.  pressure  in  each  case,  about  135  c.c.  of  ammonia,  66  c.c.  of 
carbon  dioxide,  13  c.c.  of  nitrogen,  and  1.6  c.c.  of  hydrogen.  Dewar 
found  that  i  c.c.  of  charcoal  absorbed  at  o°  and  76  cm.  pressure  4  c.c., 
at— 185°  and  76  cm.  pressure,  35  c.c.  of  hydrogen  corrected  to 
N.T.P.  Dewar  has  taken  advantage  of  this  strong  adsorptive 
power  of  charcoal  for  gases  at  low  temperatures  to  remove  the 
last  traces  of  gas  from  evacuated  flasks.  For  the  same  reason 
charcoal  is  largely  employed  as  an  absorbent  of  deleterious  gases 
in  cisterns,  etc.,  and  as  a  trap  in  sewers. 

The   adsorption   from   solution   of  substances   of  high  molecular 
weight  by  charcoal  has  already  been  referred  to.     It  can  be  shown 


FIG.  68. 


330    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

by  shaking  up  water  deeply  coloured  by  caramel  with  animal  char- 
coal and  then  filtering,  when  the  water  will  pass  through  almost 
colourless.  On  account  of  its  highly  porous  character,  animal  charcoal 
is  most  largely  used  in  this  connexion,  for  example,  to  remove 
colouring  matters  in  the  refining  of  sugar. 

Chemical  Properties  of  Carbon— When  heated  in  the  air, 
carbon  combines  with  oxygen  to  form  carbon  dioxide,  CO2,  but  the 
temperature  at  which  combination  begins  varies  enormously  with 
the  nature  of  the  carbon,  and  appears  to  be  the  higher  the  greater 
the  density.  Wood  charcoal  catches  fire  about  300°,  sugar  charcoal, 
previously  ignited  in  chlorine,  at  about  450°,  whilst  gas  carbon,  as 
already  mentioned,  can  be  raised  to  an  extremely  high  temperature 
before  igniting.  When  the  supply  of  air  is  insufficient,  carbon  mon- 
oxide, CO,  is  formed  to  some  extent. 

Carbon  combines  directly  with  hydrogen  at  1100-1200°  to  form 
marsh  gas  or  methane^  CH4.  At  the  temperature  of  the  electric  arc 
another  compound,  acetylene,  C2H2  (q.v.\  is  formed  by  direct  com- 
bination of  the  elements. 

Carbon  combines  with  most  ot  the  metals,  in  some  cases  directly, 
to  form  carbides.  The  most  important  of  these  compounds,  calcium 
carbide,  CaC2,  will  be  referred  to  at  a  later  stage  (p.  454). 

Carbon  also  combines  directly  with  certain  non-metals.  With 
sulphur  it  forms  carbon  disulphide,  CS2  (p.  341).  With  silicon  it 
combines  in  the  electric  furnace  to  form  an  extremely  hard  com- 
pound, carborundum,  CSi  (p.  367). 

Coal — As  already  indicated,  coal  is  formed  by  the  slow  decay 
of  vegetable  matter  in  the  absence  of  air.  The  process  of  coal 
formation  consists  essentially  in  the  removal  of  the  hydrogen  and 
oxygen  along  with  part  of  the  carbon  as  water,  carbon  dioxide, 
marsh  gas  or  methane  and  other  products.  The  escape  of  methane 
from  marshes  is  an  illustration  of  this  change  at  work.  Materials 
representing  all  stages  of  the  process  are  familiar  to  us,  the  order 
from  the  youngest  to  the  most  completely  changed  material  being 
as  follows :  peat,  brown  coal  or  lignite,  soft  or  bituminous  coal, 
and  hard  or  anthracite  coal.  The  average  composition  of  these 
materials,  excluding  ash  and  moisture,  as  regards  carbon,  hydrogen 
and  oxygen,  and  the  percentage  of  ash,  is  shown  in  the  accompany- 
ing table. 

Anthracite  coal  burns  with  scarcely  any  flame  or  smoke,  and 
is  largely  used  as  steam  coal.  Bituminous  coal  burns  with  a 
smoky  flame,  yields  volatile  hydrocarbons  (compounds  of  carbon 


CARBON  331 

and  hydrogen)  on   distillation,   and   is  largely   used   for   preparing 
coal  gas  and  for  ordinary  household  purposes. 


Carbon. 

Hydrogen. 

Oxygen. 

Ash. 

Wood  (cellulose)    .     . 

50 

6 

44 

[1-5] 

Peat      

60 

6 

34 

5-20 

Brown  coal    .... 

67 

5 

27 

[5-20] 

Bituminous  coal      .     . 

81 

5 

8 

fi-io] 

Anthracite      .... 

94 

3 

3 

1-2] 

The  Electric  Furnace — The  electric  furnace  as  used  by 
Moissan  is  represented  in  cross-section  in  Fig.  69.  It  consists  of 
two  blocks  of  quicklime  hollowed  out  so  as  to  make  a  reaction 
chamber  and  fitting  together  as  shown,  the  carbon  electrodes  rest- 
ing on  the  lower  block.  The  heat  is  reflected  on  the  crucible 
from  the  rounded  upper  surface  of  the  chamber.  The  tempera- 
ture attainable  in  the  electric  furnace  naturally  depends  upon  the 


FIG.  69. 

voltage  and  strength  of  current  used,  but  is  limited  by  the  degree 
of  resistance  to  temperature  of  the  furnace  material.  At  3000° 
lime  melts,  and  it  can  stand  a  temperature  of  2500°  only  for  a 
limited  period.  The  electric  furnace  is  coming  increasingly  into 
use  for  various  purposes.  In  most,  if  not  in  all  cases  the  effect 
is  purely  thermal ;  the  furnace  is  simply  a  convenient  apparatus 
for  obtaining  a  high  temperature. 

CARBON  COMPOUNDS 

General — Carbon  is  one  of  the  elements  invariably  present  in 
substances  of  animal  and  vegetable  origin,  and  it  was  formerly 
supposed  that  such  substances  could  only  be  formed  as  the  result  of 
vital  activity.  These  substances  were  therefore  classed  as  organic^ 


332     A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

in  contrast  to  substances  of  mineral  origin,  termed  inorganic. 
Although  we  now  know  that  vital  activity  is  by  no  means  in- 
dispensable to  the  production  of  carbon  compounds,  as  many  of 
them  have  been  prepared  in  the  laboratory  from  their  elements, 
it  has  nevertheless  been  found  desirable  to  retain  the  distinction 
between  organic  chemistry  and  inorganic  chemistry,  mainly  on 
account  of  the  enormous  nuHlfiHr* 'and  great  importance  .  of  the 
carbon  compounds.  The  chief  reason  for  the  existence  of  so 
many  compounds  containing  this  element  is  to  be  found  in  the 
capacity  of  carbon  atoms  to  unite  among  themselves,  forming, 
along  with  other  elements  (chiefly  hydrogen,  oxygen  and  nitrogen) 
chains  and  rings. 

Although  the  study  of  carbon  compounds  in  general  belongs  to 
organic  chemistry,  it  is  usual  to  deal  with  the  oxides  of  carbon 
and  the  carbonates  in  inorganic  chemistry  ;  in  fact,  the  com- 
pounds just  mentioned  are  not  regarded  as  organic  compounds. 
This  term  is  confined  to  compounds  containing  hydrogen  as  well 
as  carbon. 

Compounds  containing  carbon  and  hydrogen  only  are  termed  hydro- 
carbons. On  account  of  their  importance  in  connexion  with  com- 
bustion, and  for  other  reasons,  we  will  deal  very  briefly  with  three 
simple  hydrocarbons,  methane,  CH4,  ethylene,  C2H4,  and  acetylene, 
C2H2.  A  compound  of  carbon  and  nitrogen,  cyanogen,  C2N2,  and 
some  of  its  simple  derivatives  will  also  be  shortly  referred  to. 

COMPOUNDS  OF  CARBON  AND  OXYGEN 

Three  oxides  of  carbon  are  known  : 

Carbon  suboxide C3O2 

Carbon  monoxide         .        .         .        .        .        .        .CO 

Carbon  dioxide    ........     CO2. 

Carbon  dioxide  forms  a  corresponding  acid,  carbonic  acid,  H2CO3. 
Carbon  monoxide  may  also  be  regarded  as  the  anhydride  of  an  acid 
known  as  formic  acid,  HCOOH.  Carbon  suboxide  was  discovered 
so  recently  as  1906. 

Carbon  Suboxide,  C3O2>  is  obtained  by  heating  malonic  acid  with  phosphorus 
pentoxide:  CH2(COOH)2+2P2O5->C3O2+4HPO3.  It  is  a  colourless,  poisonous 
gas  with  an  acrid  odour  ;  it  can  be  obtained  as  a  colourless  liquid  which  boils  at 
7°  and  as  a  solid  melting  at  — 107°.  It  decomposes  slowly  at  room  temperature. 
Chemically  it  behaves  like  malonic  anhydride,  combining  with  water  to  form 
malonic  acid. 


CARBON  333 

CARBON  MONOXIDE,  CO 

Preparation  —  (i)   Carbon    monoxide   is   formed   when   carbon 
dioxide  is  passed  over  red-hot  carbon  : 


This  reaction  takes  place  in  a  glowing  coal  fire.  The  air  entering 
below  combines  with  the  carbon  to  form  the  dioxide,  which  on  its 
passage  upwards  through  the  glowing  fuel  is  reduced  to  the  monoxide 
in  accordance  with  the  above  equation.  When  the  monoxide  reaches 
the  surface  of  the  coals  and  comes  into  contact  with  the  air,  it  burns 
to  the  dioxide  with  a  characteristic  bluish  flame. 

The  above  reaction  is  reversible.  At  450°  the  equilibrium  mixture 
contains  very  little  (less  than  2  per  cent.)  of  carbon  monoxide,  at 
1000°,  on  the  other  hand,  the  reaction  is  practically  complete  in  the 
direction  of  the  upper  arrow. 

(2)  Carbon  monoxide  is  also  obtained  by  reducing  metallic  oxides, 
e.g.  zinc  oxide,  with  carbon  : 


(3)  By  passing  carbon  dioxide  over  heated  zinc,  aluminium  or  iron  : 

Zn  +  CO2->ZnO  +  CO. 

(4)  When  steam  is  passed  over  carbon  raised  to  a  white  heat  (over 
1000°)  a  mixture  of  carbon  monoxide  and  hydrogen  is  obtained: 


The  product,  known  as  water  gas,  is  employed  commercially  as  a 
gaseous  fuel.  With  the  addition  of  certain  hydrocarbons  rich  in  carbon, 
it  burns  with  a  luminous  flame  and  is  used  for  lighting  purposes.1 

(5)  Carbon  monoxide  is  readily  obtained  from  formic  acid  or  a 
formate  by  heating  with  concentrated  sulphuric  acid,  water  being 
abstracted  : 

HCOOH->CO  +  H2O. 

(6)  When  under  the  same  circumstances  oxalic  acid  or  an  oxalate 
is  used,  a  mixture  of  carbon  monoxide  and  dioxide  is  obtained: 

(COOH)2->CO  +  CO2  +  H2O. 

The   carbon   dioxide   can   be   removed  by  passing  the   mixed   gas 
through  a  solution  of  sodium  or  potassium  hydroxide  : 
C02  +  2NaOH->Na2CO3+H20. 

1  Producer  gas,  obtained  by  the  partial  combustion  of  coke  in  air,  is  a  mixture 
of  carbon  monoxide  and  nitrogen. 


334     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

(7)  When  the  gas  is  required  in  considerable  amount  it  is  most 
conveniently  obtained  by  heating  potassium  ferrocyanide  (p.  563) 
with  five  times  its  weight  of  concentrated  sulphuric  acid  in  a  large 
flask: 

K4Fe(CN)6  +  6H2SO4  +  6H2O->FeSO4  +  2K2SO4 
+  3(NH4)2S04  +  6CO. 

As  the  equation  shows,  the  reaction  is  very  complicated. 

Physical  Properties — Carbon  monoxide  is  a  colourless,  odour- 
less, tasteless  gas.  Under  great  pressure  at  low  temperatures  it  is 
obtained  as  a  colourless  liquid  which  boils  at  -  191°  ;  on  evaporation 
under  reduced  pressure  it  is  obtained  as  a  snow-like  solid  which 
melts  at  -208°.  The  critical  temperature  is  -  139.5°  and  the  critical 
pressure  35.5  atmospheres  (Olszewski). 

Carbon  monoxide  is  very  slightly  soluble  in  water  ;  the  coefficient 
of  absorption  is  0.0354  at  o°  and  0.0258  at  20°. 

Carbon  monoxide  is  a  very  poisonous  gas  owing  to  the  fact  that 
it  combines  with  the  haemoglobin  of  the  blood  to  form  a  bright  red 
compound,  known  as  carboxy-haemoglobin.  As  this  compound  is 
much  more  stable  than  the  compound  of  haemoglobin  with  oxygen, 
the  latter  is  unable  to  displace  carbon  monoxide  from  combination, 
and  the  blood  thus  poisoned  is  unable  to  fulfil  its  function  of  supplying 
oxygen  to  the  different  parts  of  the  body.  Carbon  monoxide  is  formed 
when  combustion  is  incomplete,  as  in  insufficiently  ventilated  stoves 
in  which  coke  or  coal  is  being  burned,  and  as  it  does  not  betray  its 
presence  by  any  characteristic  smell,  many  deaths  have  occurred 
owing  to  the  escape  of  the  gas  into  the  room.  A  small  proportion 
of  the  gas  mixed  with  air  causes  headache. 

Chemical  Properties— The  chemical  properties  of  carbon 
monoxide  are  readily  understood  by  regarding  it  as  an  unsaturated 
compound  (p.  347).  Its  graphic  formula  may  be  written  C  =  O  or 
C=O,  according  as  we  regard  the  oxygen  as  divalent  or  quadrivalent. 
In  any  case,  carbon  tends  to  function  as  a  quadrivalent  and  oxygen 
as  a  bivalent  element,  and  therefore  carbon  monoxide  readily  adds  on 
one  atom  of  a  bivalent  or  two  atoms  of  a  univalent  element  to  form 
a  saturated  compound. 

Carbon  monoxide  burns  in  air  with  a  bluish  flame  to  form  carbon 
dioxide : 

2CO  +  O2->2CO2. 

If  the  gases  are  perfectly  dry,  combination  does  not  occur. 
The  equation  indicates  that  two  volumes  of  the  monoxide  combine 


CARBON  335 

with  one  volume  of  oxygen  to  form  two  volumes  of  the  dioxide,  and 
this  can  be  confirmed  by  exploding  a  mixture  in  the  above  proportions. 

The  gas  unites  directly  with  chlorine  in  sunlight  to  form  carbonyl 
chloride  or  phosgene,  COC12,  and  with  sulphur  to  form  carbon  oxy- 
sulphide,  COS.  It  also  unites  directly  with  certain  metals,  especially 
nickel  and  iron,  forming  remarkable  compounds  known  as  carbonyls. 
The  best-known  is  nickel  carbonyl,  Ni(CO)4,  a  colourless,  mobile 
liquid,  which  boils  at  43°  under  750  mm.  pressure,  and  is  split  up  into 
its  components  on  passing  through  a  heated  tube.  Three  iron  car- 
bonyls are  known.  The  pentacarbonyl,  Fe(CO)6,  is  a  yellow  liquid 
which  boils  at  102.5°  under  atmospheric  pressure  ;  the  tetracarbonyl 
occurs  in  short,  dark-green,  lustrous,  prismatic  crystals,  and  diferronona- 
carbonyl,  Fe2(CO)9,  in  orange-red  crystals  (Dewar  and  Jones). 

Carbon  monoxide  is  readily  absorbed  by  an  ammoniacal  or  hydro- 
chloric acid  solution  of  cuprous  chloride  at  room  temperature  ;  from 
the  solutions  a  crystalline  compound  of  the  formula  (CuCl)2,CO,2H2O 
has  been  obtained.  Carbon  monoxide  does  not  combine  directly 
with  water  to  produce  formic  acid  (p.  350),  but  combines  with  solid 
alkalis  on  heating  to  form  the  corresponding  formates  : 

KOH  +  CO->HCOOK. 

On  account  of  the  readiness  with  which  it  combines  with  oxygen, 
carbon  monoxide  is  a  powerful  reducing  agent.  It  reduces  the  oxides 
of  certain  metals,  e.g.  ferric  oxide,  to  the  metal  (p.  559) : 

Fe2O3+3CO->2Fe+3CO2. 

The  composition  and  thermochemical  behaviour  of  carbon  mon- 
oxide are  referred  to  in  connexion  with  carbon  dioxide  (p.  339). 

CARBON  DIOXIDE,  CO2 

History — Carbon  dioxide  is  the  first  gas  which  was  definitely 
recognized  as  being  different  from  ordinary  air.  This  discovery  was 
made  in  the  beginning  of  the  seventeenth  century  by  van  Helmont, 
who  obtained  the  gas  by  burning  wood  and  by  the  action  of  acids  on 
chalk,  and  showed  that  it  differed  from  ordinary  air  in  not  being  a 
supporter  of  combustion.  Joseph  Black  of  Edinburgh  discovered  in 
1757  that  the  gas  was  absorbed  by  caustic  alkalis  (alkali  hydroxides), 
and  for  this  reason  called  \\-fixed  air.  Finally,  in  1781,  Lavoisier 
proved  that  "  fixed  air  "  is  a  compound  of  carbon  and  oxygen. 

Occurrence— Carbon  dioxide  is  a  regular  constituent  of  the 
atmosphere  ;  on  an  average  10,000  volumes  of  air  contain  3  volumes 


336     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

of  the  gas.  It  is  present  in  all  natural  waters  (p.  59).  Certain  mineral 
waters,  for  example  those  at  Vichy,  Selters,  and  Carlsbad,  contain 
the  gas  under  considerable  pressure,  and  it  escapes  with  effervescence 
when  the  pressure  is  reduced.  It  escapes  from  cracks  in  the  earth, 
particularly  in  volcanic  regions,  as  in  the  Grotto  del  Cane  at  Naples, 
and  is  given  off  from  active  volcanoes. 

Preparation  —  (i)  Carbon  dioxide  is  obtained  by  burning  carbon 
with  free  access  of  air  : 

C  +  O2-»CO2. 

(2)  The  carbonates  of  all  metals  except  the  normal  carbonates  of 
potassium  and  sodium  decompose  on  heating,  giving  the  correspond- 
ing oxide  and  carbon  dioxide.     In  practice  calcium  carbonate  is  gener- 
ally used  : 

CaCO3->CaO  +  CO2. 

With  proper  precautions  a  very  pure  gas  is  obtained  by  this  method. 

(3)  Carbon  dioxide  is  also  given  off  when  a  carbonate  is  decom- 
posed by  a  strong  acid,  such  as  hydrochloric  acid  : 

CaCO3+  2HCl->CaCl2  +  H2O  +  CO2. 

The  nature  of  this  reaction  will  be  readily  understood  in  the  light  of 
the  considerations  on  chemical  equilibrium  already  advanced.  As 
carbonic  acid  is  very  slightly  ionised  (p.  267)  it  is  almost  completely 
displaced  from  combination  by  hydrochloric  acid  : 

CaCO3+2HCl^CaCl2+H2CO3  (i.). 

Further,  carbonic  acid  has  a  great  tendency  to  dissociate  into  carbon 
dioxide  and  water  : 


The  equilibrium  lies  very  near  the  right-hand  side,  and  as  carbon 
dioxide  is  only  slightly  soluble  in  water  it  escapes  from  the  system, 
and  the  reaction  proceeds  according  to  equation  (i.)  in  the  direction 
of  the  upper  arrow. 

The  reader  should  himself  trace  out  the  mechanism  of  other  re- 
actions in  the  same  way. 

This  method  of  obtaining  carbon  dioxide  is  the  most  convenient 
for  laboratory  purposes,  lumps  of  marble  or  limestone  and  dilute 
hydrochloric  acid  being  used  in  a  Kipp's  apparatus  in  the  ordinary 
way  (p.  35). 

(4)  When  a  sugar,  such  as  glucose,  CaH12O6,  undergoes  fermentation 


CARBON  337 

in  the  presence  of  yeast,  ordinary  alcohol,  C2H6OH,  and  carbon  dioxide 
are  the  main  products  : 

C0H12O6->2C2H6 


(5)  Carbon  dioxide,  along  with  water  and  other  products,  is  formed 
when  organic  substances  (sugar,  petroleum,  oils,  the  materials  of 
candles,  etc.)  are  completely  burned  in  air.  or  oxygen  ;  or  when  they 
are  heated  with  compounds  which  readily  yield  oxygen,  for  example, 
copper  oxide. 

Physical  Properties  —  Carbon  dioxide  is  a  colourless  gas,  with 
a  distinct  slightly  pungent  smell.  As  its  density  referred  to  hydrogen 
is  about  22,  it  can  be  collected  by  upward  displacement  of  air.  It  can 
be  condensed  to  a  colourless  liquid,  which  boils  at  -  79°.  At  -  30° 
the  vapour  pressure  of  the  liquid  is  14  atmospheres,  at  o°  35  atmos- 
pheres, and  at  20°  58  atmospheres.  The  gas  is  sold  commercially  in 
strong  iron  cylinders,  and  is  used  in  the  preparation  of  aerated  waters 
and  for  other  purposes.  As  the  above  numbers  show,  it  exerts  con- 
siderable pressure  on  the  walls  of  the  containing  vessels.  A  certain 
proportion  of  the  commercial  product  is  obtained  by  collecting 
the  gas  escaping  from  fermentation  vats  and  compressing  it  into 
cylinders.  The  critical  temperature  of  carbon  dioxide  is  31.4°;  its 
critical  pressure  73  atmospheres.  Liquid  carbon  dioxide  is  miscible 
with  alcohol  and  with  ether,  but  does  not  mix  with  water. 

When  the  liquid  is  allowed  to  escape  from  a  small  orifice  (by 
partially  opening  the  valve  of  the  cylinder)  into  a  canvas  bag,  the 
evaporation  of  part  of  the  liquid  takes  up  so  much  heat  that  the 
remainder  solidifies  and  is  retained  by  the  bag.  Solid  carbon  dioxide 
is  a  white,  snow-like  substance  of  density  1.5,  which  under  atmos- 
pheric pressure  passes  directly  into  vapour  without  melting.  This  is 
owing  to  the  fact  that  the  temperature  at  which  it  exerts  a  vapour 
pressure  of  I  atmosphere,  -78°,  lies  below  its  melting-point  under 
these  conditions.  Under  a  pressure  of  5.1  atmospheres  it  melts  at 
-  56.4°.  Mixtures  of  solid  carbon  dioxide  with  ether  or  with  alcohol 
are  very  useful  cooling  agents  ;  in  both  cases  a  constant  tempera- 
ture of  -80°  is  obtained.  The  solid  dioxide  is  now  obtainable  in 
London  at  is.  per  Ib. 

At  room  temperature  water  dissolves  about  its  own  volume  ot  carbon 
dioxide.  The  "  absorption  coefficient  "  for  different  temperatures  is  as 
follows  :  1.713  at  o°,  1.194  at  10°,  1.019  at  I5°J  0.878  at  20°,  0.665  at 
30°,  and  0.530  at  40°-  Up  to  3  or  4  atmospheres  the  solubility  follows 
Henry's  law,  but  at  higher  pressures  the  gas  is  less  soluble  than  the 

22 


338     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

law  indicates.  This  is  shown  by  the  following  results,  obtained  by 
Wroblewski  at  12.4°  : 

Pressure  (atmos.)         I  5  10  20  30 

Solubility  1.086       5.15        9.65         17.11        23.25 

A  solution  of  this  gas  in  water,  prepared  under  a  pressure  of  3  to  4 
atmospheres,  is  known  as  soda-water.  When  a  bottle  is  opened 
the  excess  of  gas  above  that  corresponding  with  the  solubility  at 
atmospheric  pressure  does  not  all  escape  ;  in  other  words,  the 
solution  remains  supersaturated.  The  gas  is  completely  expelled 
by  boiling. 

The  heat  of  combustion  of  amorphous  carbon  (12  grams)  in  oxygen 
amounts  to  96,980  cal.  ;  for  diamond  the  value  is  about  94,300  cal. 
The  heat  of  combustion  of  carbon  monoxide  to  the  dioxide  (at  con- 
stant pressure)  is  68,000  cal.  It  follows,  according  to  Hess's  law, 
that  in  the  reaction 


96,980—68,000  =  28,980  cal.  are  given  out,  so  that  much  less  heat  is 
given  out  when  carbon  monoxide  is  formed  from  carbon  than  when  the 
latter  unites  with  a  second  atom  of  oxygen  to  form  the  dioxide.  This 
may  perhaps  be  explained  by  supposing  that  the  molecule  of  carbon 
is  very  complex,  and  that  part  of  the  heat  is  taken  up  in  simplifying 
the  molecule. 

Chemical  Properties  —  Carbon  dioxide  is  an  extremely  stable 
substance.  Even  at  1300°  its  dissociation,  represented  by  the  equa- 
tion 2CO2^2CO  +  O2,  is  only  0.06  per  cent.  It  is  decomposed  to  a 
considerable  extent  by  passing  electric  sparks  through  it,  the  degree 
of  decomposition  depending  upon  the  pressure  and  other  factors. 

It  does  not  support  combustion  under  ordinary  conditions,  but 
certain  substances  having  a  great  affinity  for  oxygen,  such  as  the 
alkali  metals  and  magnesium,  burn  in  the  gas.  If  a  piece  of 
ignited  magnesium  ribbon  is  immersed  in  carbon  dioxide,  it  continues 
to  burn  and  carbon  is  set  free.  With  metals  having  less  affinity  for 
oxygen,  such  as  aluminium,  zinc,  and  iron,  the  reduction  proceeds 
only  to  carbon  monoxide.  Carbon  dioxide  is  used  as  a  fire  extin- 
guisher, as  it  prevents  access  of  oxygen,  and  does  not  itself  under 
these  conditions  support  combustion. 

The  solution  of  carbon  dioxide  in  water  has  a  slightly  acid  taste 
and  turns  litmus  wine-red,  indicating  the  presence  of  an  acid.  The 
acid  has  not  been  isolated,  but  from  analogy  with  its  salts,  the  car- 


CARBON  339 

bonates,  we  assume  that  its  formula  is  H2CO3.     It  is  a  very  weak  * 
dibasic  acid,  being  ionised  mainly  according  to  the  equation 


The  second  stage  of  the  ionisation,  HCO3/^tH'  +  CO3//,  is  so  slight  as 
to  be  practically  negligible. 

Whether  carbon  dioxide  in  aqueous  solution  is  partly  present  as 
such  and  partly  in  the  hydrated  form,  or  whether  it  is  all  present  in 
the  hydrated  form,  is  not  definitely  known  ;  but  the  former  alternative 
is  the  more  probable. 

Animals  brought  into  an  atmosphere  of  carbon  dioxide  die  on 
account  of  the  absence  of  oxygen.  Apart  from  this,  however,  the  gas 
has  a  direct  anaesthetic  action  ;  and  air  containing  a  proportion  of 
the  gas  much  above  the  average  amount  should  not  be  inhaled. 

Composition  of  Carbon  Dioxide  and  Carbon  Mon- 
oxide —  Atomic  Weight  of  Carbon—  It  can  be  shown,  by 
means  of  the  apparatus  described  in  connexion  with  sulphur  dioxide, 
that  carbon  dioxide  contains  its  own  volume  of  oxygen.  It  follows 
that  the  molecule  of  the  dioxide  contains  one  molecule  or  two  atoms 
of  oxygen,  and  that  its  formula  is  Q*O2.  Direct  determination  of  its 
composition  shows  that  it  contains  12  parts  of  carbon  to  32  parts  of 
oxygen,  and  as  no  gaseous  compound  is  known  whose  molecule  con- 
tains less  than  12  parts  of  carbon,  the  atomic  weight  of  this  element 
must  be  12  and  the  formula  for  carbon  dioxide  CO2.  These  considera- 
tions are  confirmed  by  the  observation  that  the  den'sity  of  the  gas 
is  22  ;  the  molecular  weight  is  therefore  44. 

As  above  indicated,  the  accurate  atomic  weight  of  carbon  can  be 
found  by  determining  the  exact  ratio  in  which  it  combines  with  oxygen 
to  form  carbon  dioxide.  This  was  done  in  a  classical  investigation  by 
Dumas  and  Stas  (1844).  A  platinum  boat  containing  either  diamond 
or  graphite  was  carefully  weighed  and  placed  in  a  porcelain  tube, 
which  was  raised  to  a  red  heat  while  carefully  purified  oxygen  was 
passed  through  it.  The  carbon  dioxide  formed  was  completely 
absorbed  in  bulbs  containing  potassium  hydroxide,  which  were 
weighed  before  and  after  the  experiment,  while  the  amount  of  carbon 
burned  was  obtained  by  weighing  the  boat  and  its  contents  after  the 
experiment.  The  mean  of  a  number  of  experiments  was  as  nearly  as 
possible  :  Carbon  :  oxygen  =  3  :8,  whence  it  follows  that  (for  oxygen 
=  32)  the  atomic  weight  of  carbon  is  12.00.  The  most  trustworthy 
later  determinations  of  this  constant  have  fully  confirmed  the  accuracy 
of  this  figure. 


340     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

As  carbon  monoxide  combines  with  half  its  volume  of  oxygen  to 
form  carbon  dioxide,  whereas  the  latter,  as  shown  above,  contains  its 
own  volume  of  oxygen,  it  follows  that  carbon  monoxide  contains  half 
its  volume  of  oxygen.  Hence,  according  to  Avogadro's  hypothesis, 
each  molecule  of  carbon  monoxide  contains  one  atom  of  oxygen.  As 
the  molecular  weight  of  the  compound,  determined  from  its  vapour 
density,  is  28,  it  follows  at  once  that  its  formula  is  CO. 

Carbonates  —  Although,  as  we  have  seen,  carbonic  acid,  H2CO3, 
is  unstable,  the  corresponding  salts,  the  carbonates,  are  in  most  cases 
stable.  As  carbonic  acid  is  dibasic,  it  forms  with  univalent  metals 
two  series  of  salts  —  acid  salts,  e.g.  NaHCO3,  and  normal  salts,  e.g^ 
Na2C03. 

When  carbon  dioxide  is  passed  in  excess  through  a  solution  of 
sodium  hydroxide,  sodium  hydrogen  carbonate  (sodium  bicarbon- 
ate) is  formed  : 

NaOH  +  CO2->NaHCO3. 

The  salt  is,  of  course,  ionised  in  solution  according  to  the  equation 
NaHCO3^±Na'  +  HCO3',  and  as  the  HCO3'  ion  splits  off  practically 
no  H-  ions,  the  solution  is  very  nearly  neutral.  When  sodium 
bicarbonate  is  mixed  with  an  equivalent  of  sodium  hydroxide,  sodium 
carbonate,  Na2CO3,  is  formed  : 


The  solution'  of  normal  sodium  carbonate  in  water  is  alkaline, 
because  it  is  partially  hydrolyzed,  as  represented  by  the  lower  arrow 
in  the  above  equation.  This  hydrolysis  is  connected  with  the 
exceeding  weakness  of  carbonic  acid  as  a  dibasic  acid  (cf.  sulphites, 
p.  302). 

When  carbon  dioxide  is  passed  into  a  solution  of  calcium  hy- 
droxide, calcium  carbonate  is  at  first  precipitated  according  to 
equation  (i)  : 

(I)  Ca(OH)2  +  C02-»CaC03+H20 
(2) 


but  on  continuing  the  passage  of  the  gas  a  clear  solution  is  finally 
obtained.  Further  investigation  shows  that  the  solution  contains  an 
acid  calcium  carbonate,  Ca(HCO3)2,  derived  from  two  molecules  of 
carbonic  acid  by  displacing  two  hydrogen  atoms  by  an  atom  of 
calcium.  When  the  solution  is  boiled,  the  reaction  is  completely 
reversed  in  the  direction  of  the  lower  arrow,  the  calcium  carbonate 


CARBON  34i 

being  reprecipitated.  The  calcium  carbonate  present  in  solution  in 
natural  waters  occurs  as  the  soluble  acid  carbonate,  and  is  partly 
responsible  for  the  hardness  of  such  waters.  The  insoluble  carbonates 
of  certain  other  metals,  e.g.  magnesium,  are  also  held  in  solution  by 
excess  of  carbon  dioxide. 

The  carbonates  of  the  alkali  metals  are  readily  soluble  in  water, 
those  of  the  other  metals  are  nearly  all  insoluble,  and  are  therefore 
prepared  by  double  decomposition.  The  normal  alkali  carbonates 
are  stable  even  at  high  temperatures  ;  the  bicarbonates  lose  carbon 
dioxide  on  heating,  and  the  normal  carbonates  are  left  : 

2NaHCO3->Na2CO3+  H2O  +  CO2. 

The  carbonates  of  most  other  metals  readily  yield  the  oxide  and 
carbon  dioxide  on  heating. 

Percarbonic  Acid,  H2C2O6  (probably  HOCO'OOCOOH)  is  obtained  as  a 
salt  by  electrolysis  of  a  concentrated  solution  of  an  alkali  carbonate  or  bicarbonate. 
The  free  acid  is  unstable,  decomposing  into  carbon  dioxide  and  hydrogen 
peroxide. 

CARBON  BISULPHIDE,  CS2 

Preparation  —  Carbon  di  sulphide  is  obtained  when  the  vapour 
of  sulphur  is  passed  over  heated  carbon  (coke  or  charcoal).  The 
method  now  most  largely  used  on  the  commercial  scale  is  due  to  Taylor. 
The  heating  is  effected  electrically.  Coke  is  fed  in  between  the 
electrodes,  and  sulphur  is  introduced  just  below.  The  heat  volatilizes 
the  sulphur,  which  passes  over  the  glowing  coke  forming  vapours  of 
the  disulphide,  which  are  led  off  and  condensed.  The  process  is 
a  continuous  one.  The  disulphide  can  be  purified  by  shaking  with 
metallic  mercury  and  mercuric  sulphate  and  then  distilling. 

Properties  —  Carbon  disulphide  is  a  colourless  liquid  with  an 
agreeable  aromatic  odour  when  pure,  but  the  commercial  article  has 
a  very  disagreeable  odour,  owing  to  impurities.  Its  density  is  1.264 
at  20°  ;  it  boils  at  46.3°.  It  burns  in  air  or  oxygen  with  a  blue  flame 
to  carbon  dioxide  and  sulphur  dioxide  : 


and  its  vapour  forms  a  highly  explosive  mixture  with  air. 
The  equilibrium  between  carbon  and  sulphur  — 


at  ordinary  temperatures  lies  very  near  the  right  side  of  the  equation, 
so  that  the  disulphide  is  an  unstable  substance.     The  shock  due  to 


342     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  explosion  of  fulminating  mercury  causes  it  to  split  up  explosively 
into  its  components.  The  rate  of  the  reaction  represented  by  the 
upper  arrow  is,  of  course,  increased  by  raising  the  temperature.  The 
disulphide  is  an  endothermic  compound,  the  thermochemical  equation 
representing  its  formation  being  as  follows  :  — 

(liquid)  -22,300  cal., 


so  that  its  tendency  to  decompose  into  its  elements  is  readily  under- 
stood (p.  147). 

Carbon  disulphide  is  an  excellent  solvent  for  fats,  oils,  bromine, 
iodine,  sulphur,  phosphorus,  and  other  substances,  and  is  employed 
commercially  in  vulcanizing  rubber,  for  extracting  essential  oils  from 
plants,  etc. 

When  potassium  sulphide  is  heated  with  carbon  disulphide, 
potassium  thiocarbonate,  K2CS3,  is  obtained.  From  the  latter  sub- 
stance by  the  action  of  dilute  acids,  thiocarbonic  acid,  H2CS3,  is 
liberated  as  a  yellow,  oily  liquid,  which  decomposes  readily  into 
carbon  disulphide  and  hydrogen  sulphide.  The  close  analogy  of 
these  compounds,  CS2,  K2CS3,  and  H2CS3,  with  the  oxygen  com- 
pounds, CO2,  K2CO3,  and  H2CO3,  is  evident,  and  form  a  further 
illustration  of  the  statement  (p.  322)  that  oxygen  and  sulphur  belong 
to  the  same  family  of  elements. 

Carbon  Oxy  sulphide,  COS,  prepared  by  passing  carbon 
monoxide  and  sulphur  vapour  through  a  tube  at  a  moderate  tem- 
perature, is  a  colourless,  odourless  gas,  which  can  be  converted  into 
a  liquid  boiling  at  —  50°.  It  burns  in  the  air  with  a  blue  flame  to 
carbon  dioxide  and  sulphur  dioxide.  It  is  readily  soluble  in  water; 
but  is  slowly  decomposed  in  solution  with  formation  of  carbon 
dioxide  and  hydrogen  sulphide. 

Carbonyl  Chloride  and  Urea.  Carbonyl  chloride  or  phos- 
gene, COC12,  is  obtained  by  mixing  equal  volumes  of  carbon  monoxide 
and  chlorine  and  exposing  to  sunlight  (p.  335),  or  by  passing  the  mixed 
gases  over  heated  charcoal  : 

CO  +  C12->COC12. 

At  the  ordinary  temperature  carbonyl  chloride  is  a  colourless  gas  with 
a  suffocating  odour  ;  it  can  readily  be  condensed  to  a  liquid  which 
boils  at  8°.  When  treated  with  water  it  is  decomposed  into  carbon 
dioxide  and  hydrochloric  acid  : 


CARBON  343 


When  treated  with  ammonia  it  forms  urea  : 
Cl  /NH2 


O:C(       +4NHt-*O:C<  +2NH4C1. 

\C1 


The  urea  can  be  separated  from  the  ammonium  chloride  simul- 
taneously formed  by  taking  advantage  of  its  solubility  in  alcohol, 
from  which  it  separates  in  colourless  crystals. 

It  is  clear  from  its  formula  that  urea  may  be  regarded  as  being 
derived  from  carbonic  acid,  CO(OH)2,  by  displacement  of  the  two 
hydroxyl  groups  by  NH2  or  amido  groups.  Substances  derived  from 
acids  in  this  way  are  called  amides.  Urea  is  therefore  the  amide  of 
carbonic,  acid,  and  for  this  reason  is  also  called  carbamide.  It  can 
easily  be  hydrolyzed  to  carbon  dioxide  and  ammonia  : 

/NH2 
CO<  +H2O->CO2  +  2NH3. 

\NH2 

Urea  is  formed  in  the  animal  body  as  the  result  of  the  decomposition 
of  nitrogenous  compounds,  and  is  excreted  in  the  urine,  from  which 
it  can  be  obtained  by  evaporation.  It  was  formerly  considered  that 
urea  and  other  "  organic  "  compounds  could  only  be  obtained  as  the 
result  of  vital  activity  (p.  331).  In  1828,  however,  Wohler  made  the 
important  discovery  that  urea  can  be  obtained  by  heating  ammonium 
cyanate  (see  below)  :  , 

NH4CNO->CO(NH2)2, 

a  compound  which  can  be  prepared  in  the  laboratory  from  its 
elements.  As  a  result  of  this  and  other  discoveries  it  came  to  be 
recognized  that  there  is  no  essential  difference  between  the  so-called 
"inorganic"  and  "organic"  compounds. 

Cyanogen  and  allied  Compounds—  Cyanogen,  C2N2, 
appears  to  be  formed  by  direct  combination  of  its  elements  when  the 
electric  arc  is  passed  between  carbon  poles  in  an  atmosphere  of 
nitrogen,  but  it  cannot  be  detected  in  the  gases  drawn  from  the  arc 
chamber. 

It  is  readily  obtained  by  heating  dry  mercuric  cyanide  ' 

Hg(CN)2->Hg  +  C2N2, 

and  by  the  action  of  potassium  cyanide  on  a  solution  of  copper 
sulphate  : 


344     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

Cyanogen  is  a  colourless  gas,  which  is  fairly  soluble  in  water,  and 
burns  in  air  with  a  lavender-coloured  flame  to  carbon  dioxide  and 
nitrogen  : 

C2N2  +  2O2->2CO2  +  N2. 

Like  carbon  disulphide  and  acetylene  (g.v.)  it  is  an  endothermic 
compound.  In  its  chemical  properties  cyanogen  shows  a  remarkable 
resemblance  to  the  halogens.  It  combines  with  hydrogen  at  high 
temperatures  to  form  a  monobasic  acid,  hydrocyanic  acid,  HCN. 
When  passed  into  an  alkali  hydroxide,  a  mixture  of  alkali  cyanide 
and  cyanate  is  obtained  (cf.  chlorine,  p.  179)  : 

C2N2+2KOH->KCN  +  KCNO  +  H2O.          % 

Hydrocyanic  Acid,  HCN,  is  prepared  by  direct  combination 
of  cyanogen  and  hydrogen,  or  carbon,  nitrogen  and  hydrogen,  in  the 
electric  arc,  but  most  readily  by  heating  potassium  cyanide  with 
sulphuric  acid  and  collecting  the  distillate  in  a  cooled  receiver. 

Pure  hydrocyanic  acid  is  a  colourless  liquid,  with  an  odour  like 
that  of  bitter  almonds  ;  it  boils  at  26.5°.  It  acts  as  an  extremely  weak 
monobasic  acid,  so  that  the  corresponding  salts,  the  cyanides,  are 
hydrolyzed  to  a  considerable  extent  in  aqueous  solution. 

Cyanates — When  potassium  cyanide  is  fused  with  lead  oxide, 
the  latter  is  reduced  to  metallic  lead,  and  potassium  cyanate,  KCNO, 
is  formed  : 

KCN  +  PbO->KCNO  +  Pb. 

As  has  already  been  mentioned,  the  cyanate,  mixed  with  cyanide, 
is  obtained  by  the  action  of  cyanogen  on  potassium  hydroxide. 

The  corresponding  ammonium  salt,  NH4CNO,  a  colourless  crystal- 
line substance,  changes  on  heating  into  urea,  as  already  mentioned. 

Thiocyanates — As  the  name  indicates,  the  salts  are  derived 
from  the  cyanates  by  displacement  of  the  oxygen  by  sulphur.  When 
potassium  cyanide  is  evaporated  down  with  ammonium  sulphide, 
the  free  sulphur  present  in  the  latter  (p.  421)  converts  the  cyanide 
into  thiocyanate.  The  same  change  can  be  effected  by  boiling  an 
aqueous  solution  of  potassium  cyanide  with  sulphur  : 

KCN  +  S-»KCNS. 

As  already  mentioned  (p.  168,  cf.  also  p.  564),  thiocyanates  give  a 
deep  red  colour  with  ferric  salts,  owing  to  the  formation  of  ferric 
thiocyanate. 

t 


CARBON 


345 


SOME  SIMPLE  ORGANIC  COMPOUNDS 

For  the  proper  understanding  of  certain  phenomena  such  as  the 
nature  of  flame  and  the  carbon  cycle  in  nature,  it  will  be  necessary 
to  refer  to  certain  compounds  which  belong  to  the  organic  division 
of  the  subject.  These  compounds  are  therefore  very  briefly  con- 
sidered here  ;  for  a  full  discussion  of  their  properties  books  on 
organic  chemistry  must  be  consulted. 

It  has  already  been  stated  that  compounds  which  contain  carbon 
and  hydrogen  only  are  called  hydrocarbons.  Four  hydrocarbons, 
methane,  ethane,  ethylene,  and  acetylene  are  of  importance  for  our 
present  purpose. 

METHANE  OR  MARSH  GAS,  CH4 

Occurrence  —  The  gas  escaping  in  bubbles  from  marshes,  which 
results  from  the  slow  decay  of  organic  matter,  contains  a  large  pro- 
portion (usually  more  than  80  per  cent.)  of  methane,  hence  the  name 
marsh  gas.  It  is  also  produced  in  large  amount  during  the  forma- 
tion of  the  coal  measures,  and  is  enclosed,  often  under  considerable 
pressure,  in  the  coal.  It  escapes  during  the  working  of  the  coal 
seams,  and  is  the  chief  constituent  of  the  fire-damp  which  forms  a 
dangerously  explosive  mixture  with  air.  It  is  also  the  main  con- 
stituent of  the  "  natural  gas  "  which  escapes  from  the  earth  in  coal 
and  petroleum  districts. 

Modes  of  Formation  —  (i)  Methane  is  obtained  by  direct 
combination  of  its  elements  when  hydrogen  is  passed  over  sugar 
charcoal  heated  in  a  porcelain  tube  to  1100-1150°. 

(2)  By  passing  carbon  disulphicle  vapour  and  hydrogen  sulphide 
over  heated  copper  : 


(3)  By  heating  dry  sodium  acetate  with  sodium  hydroxide  : 
CH3COONa  +  NaOH->CH4+Na2CO3. 

Properties  —  Methane  is  a  colourless,  odourless,  tasteless  gas. 
It  burns  in  excess  of  air  or  oxygen  with  a  slightly  luminous  flame  to 
carbon  dioxide  and  water  : 


and  forms  a  highly  explosive  mixture  with  air.    On  account  of  its  con- 
nexion with  explosions  in  coal  mines  the  properties  of  this  mixture 


346     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

have  been  frequently  investigated  from  this  point  of  view.  It  appears 
that  a  mixture  of  air  and  methane  containing  as  little  as  2  per  cent. 
of  the  latter  is  dangerous. 

Chemically  methane  is  a  very  inactive  substance.  It  is,  however, 
acted  on  by  chlorine  and  bromine.  In  diffused  daylight  the  hydrogen 
atoms  are  slowly  displaced  in  successive  stages  by  the  former  gas, 
according  to  the  equations  — 


CH3C1  +  Cljr>CH2CI2  +  HC1, 

and  so  on  till  finally  all  the  hydrogen  is  displaced.  Reactions  of  this 
type  are  called  substitution  reactions,  so  that  the  compound  CH3C1  is 
a  substitution  product  of  methane.  Such  reactions  are  very  frequent 
in  organic  chemistry. 

ETHANE,  C2H6 

Preparation  —  Ethane  can  readily  be  obtained  from  the  sub- 
stitution product  of  methane,  CH3Br,  by  treatment  with  metallic 
sodium  : 


Properties  —  Ethane  is  a  colourless  gas,  which  resembles 
methane  very  closely  in  all  its  physical  and  chemical  properties. 
From  its  mode  of  formation  it  may  be  assumed  that  its  formula  is 
H3C-CH3,  the  carbon  atoms  being  quadrivalent  and  the  hydrogen 
atoms  univalent.  The  univalent  group  CH3,  which  occurs  very 
frequently  in  organic  compounds,  is  called  the  methyl  group,  the 
group  CH3'CH2  or  C2H5  is  called  the  ethyl  group. 

The  compound  C2H5C1,  which  is  a  substitution  product  of  ethane, 
is  called  ethyl  chloride,  as  it  contains  the  C2H6  or  ethyl  group.  When 
ethyl  chloride  is  heated  with  alkali  hydroxide,  the  chlorine  is  displaced 
by  the  univalent  OH  group,  forming  the  compound  C2H6OH,  which 
is  ordinary  alcohol  : 

C3H6C1  +  KOH->C2H6OH  +  KC1. 

The  term  alcohol  is  a  general  one,  applied  to  compounds  which  are 
derived  from  hydrocarbons  by  the  displacement  of  one  or  more  atoms 
of  hydrogen  by  a  corresponding  number  of  hydroxyl  groups,  and  to 
distinguish  it  from  other  alcohols  the  compound  C2H6OH  is  known 
as  ethyl  alcohol. 


CARBON  347 


ETHYLENE,  C2H4 

Preparation — (i)  This  hydrocarbon  is  obtained  by  heating 
ethyl  alcohol  with  concentrated  sulphuric  acid  to  150°  : 

C2H5OH->C2H4  +  H20. 

The  gas  is  purified  by  bubbling  it  through  a  solution  of  sodium 
hydroxide,  and  may  be  collected  over  water.  From  the  above  equa- 
tion it  would  appear  that  the  process  is  one  of  simple  dehydration,  but 
in  reality  it  is  probably  more  complicated. 

(2)  A  much  purer  product  is  obtained  by  heating  alcohol  with  con- 
centrated phosphoric  acid  at  200  to  220°.  The  final  result  is  dehydra- 
tion of  alcohol,  as  represented  by  the  above  equation. 

Properties — Ethylene  is  a  colourless  gas  with  an  agreeable 
characteristic  odour.  It  can  be  condensed  to  a  colourless  liquid  which 
boils  at  — 103.4°.  It  burns  in  air  with  a  luminous  flame  to  carbon 
dioxide  and  water,  and  is  one  of  the  constituents  to  which  ordinary 
coal  gas  owes  its  luminosity. 

Ethylene  combines  directly  with  two  atoms  of  hydrogen,  when  the 
gases  are  passed  through  a  heated  tube,  to  form  ethane,  and  it  also 
combines  directly  with  the  halogens  to  form  halogen  derivatives  of 
ethane,  e.g.  C2H4Br2.  In  these  respects  ethylene  behaves  as  an 
unsaturated  compound.  As  regards  its  formula,  only  two  hydrogen 
atoms  are  available  for  each  atom  of  carbon,  and  in  order  to  show 
carbon  as  quadrivalent,  we  may  assume  that  the  carbon  atoms  are 
attached  by  two  valencies  instead  of  by  one,  as  in  ethane,  thus  : 

H\  /H 

/C  =  C\       .     Such  compounds,  however,  in  which  carbon  atoms 

H/  \H 

are  joined  by  more  than  one  bond,  have  a  great  tendency  to  combine 
with  elements  or  compounds  with  formation  of  substances  which  con- 
tain carbon  atoms  joined  by  single  bonds  only.  This  is  the  explana- 
tion of  the  tendency  of  ethylene  to  add  on  hydrogen,  the  halogens, 
etc.,  as  mentioned  above. 

Organic  compounds  which  contain  only  carbon  atoms  joined  by 
single  valencies  are  said  to  be  saturated;  they  have  no  additive 
character.  Methane  and  ethane  belong  to  this  class.  Compounds 
containing  carbon  atoms  united  by  two  or  three  bonds  are  said  to  be 
unsaturated;  they  are  characterized  by  their  additive  character. 
Ethylene  is  an  excellent  example  of  an  unsaturated  compound. 


348     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

ACETYLENE,  C2H2 

Preparation — (i)  Acetylene  is  formed  by  direct  combination  of 
its  elements  in  the  electric  arc  : 

C2+H2->C2H2. 

(2)  It  is  obtained  by  the  action  of  water  or  dilute  acids  on  calcium 
carbide  (p.  454) : 

CaC2  +  2H2O->C2H2  +  Ca(OH)2. 

(3)  Acetylene  is  formed  when  coal  gas  burns  in  an  insufficient 
supply  of  air  (cf.  p.  359). 

Properties — Acetylene  is  a  colourless  gas  with  a  disagreeable 
odour,  and  when  inhaled  in  considerable  amount  is  poisonous.  It  is 
soluble  in  its  own  volume  of  water  at  room  temperature.  It  is  a 
highly  endothermic  compound,  the  thermochemical  equation  ex- 
pressing its  formation  being  as  follows : — 

2C  +  H2->C2H2  -  53,200  cal. 

We  may  therefore  expect  it  to  be  an  unstable  substance,  and  as  a 
matter  of  fact,  when  under  increased  pressure,  a  shock  such  as  that 
produced  by  exploding  mercury  fulminate  decomposes  it  into  its 
elements  with  a  violent  explosion.  It  burns  under  ordinary  circum- 
stances with  a  very  luminous  but  somewhat  smoky  flame.  With  a 
special  form  of  jet,  however,  so  arranged  as  to  secure  free  access  of 
oxygen,  it  gives  a  very  white  luminous  flame,  and  is  therefore  largely 
used  for  illuminating  purposes.  It  is  one  of  the  illuminating  con- 
stituents of  coal  gas  (see  below). 

The  graphic  formula  for  acetylene  is  HC=CH.  It  is  still  more 
unsaturated  than  ethylene,  and  combines  directly  with  hydrogen, 
the  halogens,  etc.  Its  most  characteristic  property  is  the  formation 
of  a  red  precipitate  of  copper  acetylide,  C2Cu2,  when  the  gas  is 
bubbled  through  an  ammoniacal  solution  of  a  cuprous  salt.  Copper 
acetylide  is  endothermic,  like  acetylene  itself,  and  highly  explosive. 
Yellowish-white  silver  acetylide,  C2Ag2,  is  precipitated  when  acetylene 
is  passed  through  an  aqueous  solution  of  a  silver  salt. 

When  acetylene  is  passed  through  a  red-hot  tube,  it  is  partially 
polymerized  to  benzene,  C6H6,  a  hydrocarbon  also  present  in  coal  gas. 

Coal  Gas — As  the  chief  constituents  of  coal  gas  have  just  been 
described,  it  will  be  convenient  to  deal  here  with  its  preparation  and 
composition.  When  coal  is  subjected  to  dry  distillation  in  closed  iron 
retorts,  the  chief  products  are — (i)  coal  gas  ;  (2)  coal  tar  and  (3) 


CARBON  34g 

ammoniacal  liquor,  which  are  condensed  in  the  tubes  leading  from 
the  retorts  ;  (4)  coke,  which  remains  behind  in  the  retort.  The  coal 
gas  is  freed  from  a  number  of  impurities,  sulphur  compounds,  which 
are  particularly  deleterious,  being  removed  by  passing  it  over  a 
mixture  of  chalk  and  ferric  oxide,  and  is  collected  over  water.  It  is  a 
mixture  of  gases,  which  may  be  divided  into  three  groups  :  Illuminat- 
ing constituents,  heating  constituents  and  impurities.  The  more  im- 
portant members  of  the  three  classes  are  as  follows  : 

Illuminating)  Ethylene,  acetylene,  benzene  and  other  unsaturated 
constituents  f  hydrocarbons. 

Heating    1 


constitu  nts      -)  methane,  carbon  monoxide. 
Impurities         Carbon  dioxide,  nitrogen,  (hydrogen  sulphide). 

The  composition  of  coal  gas  depends  greatly  on  the  nature  of  the  coal 
used  and  on  the  temperature  to  which  it  is  heated.  The  average 
composition  in  volumes  percent,  is  shown  in  the  accompanying  table. 

Hydrogen         .         .        49  Illuminants       .        .        4 

Methane  ...         34  Nitrogen  ...        4 

Carbon  monoxide     .  8  Carbon  dioxide  .         .         I 

Good  gas  should  be  practically  free  from  sulphur  compounds. 

As  coal  gas  is  largely  used  as  a  fuel  as  well  as  for  lighting  pur- 
poses, its  value  for  heating  purposes,  so-called  calorific  power,  is  in 
many  respects  as  important  as  its  illuminating  power.  The  calorific 
power  is  determined  by  burning  a  definite  quantity  in  a  calorimeter, 
and  passing  the  products  of  combustion  through  tubes  immersed  in 
water,  so  that  all  its  heat  is  given  up  to  the  water  (p.  362). 

Ethyl  Alcohol  —  Preparation  —  (  I  )  As  already  mentioned,  ethyl 
alcohol  is  obtained  by  the  action  of  potassium  hydroxide  on  the 
ethyl  halides  : 

C2H6Br  +  KOH->C2H5OH  +  KBr. 

(2)  It  is,  however,  always  prepared  on  the  large  scale  by  fermenta- 
tion of  sugar.  To  an  aqueous  solution  of  sugar  (obtained  from 
various  sources)  yeast  is  added,  and  the  mixture  kept  about  25°. 
The  sugar  is  slowly  decomposed  into  alcohol  and  carbon  dioxide  : 

C6H12Ofl->2C2H6OH  +*2CO2. 

As  a  result  of  recent  investigations,  the  action  of  the  yeast  in  this 
process  is  now  fairly  well  understood.  It  gives  rise  to  an  organic 


350     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

catalytic  agent,  zymase,  under  the  influence  of  which  the  reaction 
represented  by  the  above  equation  takes  place.  It  is  of  interest 
to  note  that  at  present  no  other  method  of  effecting  the  above 
change  is  known. 

Properties — Ethyl  alcohol  is  a  colourless  liquid  with  a  characteristic 
odour ;  it  boils  at  78°.  It  is  the  active  constituent  of  beer,  wines, 
whisky,  brandy  and  other  beverages  prepared  by  "alcoholic" 
fermentation. 

When  acted  upon  by  oxidizing  agents,  ethyl  alcohol  loses  two 
atoms  of  hydrogen,  and  a  compound  CH3CHO,  belonging  to  the 
class  of  aldehydes,  is  obtained.  On  further  oxidation,  an  atom  of 
oxygen  is  taken  up  and  acetic  acid,  CH3COOH,  is  formed. 

Aldehydes  and  Acids— The  steps  in  the  oxidation  of  ethyl 
alcohol,  referred  to  in  the  previous  paragraph,  may  be  represented 
as  follows  : — 

CH3'CH2OH     ->     CH3'CHO     ->    CH3'COOH 
Ethyl  alcohol        Acetic  aldehyde         Acetic  acid. 

It  is  shown  in  books  on  organic  chemistry  that  all  aldehydes  contain 
the  CHO  group  of  elements,  and  the  general  formula  for  alde- 

^O 

hydes  may  therefore  be  written  R-C— H  where  R  is  a  univalent 
atom  or  group,  such  as  the  CH3  or  methyl  group.  Aldehydes 
have  the  property  of  taking  up  an  atom  of  oxygen  and 
forming  compounds  of  the  general  formula  R'COOH,  where 
R  has  the  same  meaning  as  above.  Substances  of  this  type 
are  acids,  and  just  as  all  nitrates  contain  the  NO3  group,  and  all 
aldehydes  the  CHO  group,  so  all  organic  acids  contain  the  uni- 

^O 

valent  —  C— O  —  H  or  Carboxyl  group.  An  acid  containing  one 
such  group  is  a  monobasic  acid,  e.g.  CH3COOH,  and  the  hydro- 
gen of  the  carboxyl  group  can  be  displaced  by  metals  with  formation 
of  salts,  e.g.  CH3COONa,  sodium  acetate. 

From  considerations  of  space,  the  properties  of  alcohols,  aldehydes 
and  organic  acids  cannot  be  described  more  fully  here.  It  is  of 
interest  to  mention,  however,  that  when  the  univalent  group  R, 
referred  to  above,  is  a  hydrogen  atom  instead  of  the  methyl  group, 
three  very  important  compounds  are  represented  : 

H'CH2OH  orCH3OH»       H'CHO  or  CH2O  H'COOH 

Methyl  alcohol  Formic  aldehyde          Formic  acid. 

Formic  acid  has  already  been  referred  to  in  connexion  with  carbon 


CARBON  351 

monoxide  (p.  335).  It  will  be  necessary  to  refer  to  formaldehyde  in  the 
next  section.  Methyl  alcohol  and  acetic  acid  are  two  of  the  chief 
volatile  products  formed  by  the  destructive  distillation  of  wood 

(P-  327)- 

The  Carbon  Cycle  in  Nature — It  is  a  familiar  fact  that  the 
vital  activity  of  animals  and  plants  is  attended  by  the  oxidation  of 
complex  substances  by  the  oxygen  of  the  air,  which  is  taken  in  by  the 
lungs  and  conveyed  to  the  different  parts  of  the  body  by  the  red 
colouring  matter  of  the  blood.  The  ultimate  oxidation  products,  as 
far  as  carbon  and  hydrogen  are  concerned,  are  carbon  dioxide  and 
water,  the  former  of  which  is  contained  in  expired  air. 

At  the  same  time  a  process  almost  exactly  the  converse  of  the 
above  takes  place  in  the  green  parts  of  plants  under  the  influence 
of  sunlight,  as  a  result  of  which  water  and  carbon  dioxide,  the 
latter  taken  in  from  the  atmosphere,  are  converted  into  complex 
substances  relatively  rich  in  carbon  and,  hydrogen,  and  an  amount 
of  oxygen  equivalent  to  that  of  the  carbon  dioxide  taken  in  is 
set  free.  If  we  assume  for  simplicity  that  the  substance  first, 
formed  in  the  plants  is  a  simple  sugar,  for  example  glucose,  the 
above  changes  may  be  represented  by  the  following  equation  : 
6CO2+6H2O->C6H12O6  +  6O2  (i). 

These  compounds,  relatively  poor  in  oxygen,  are  built  up  into 
starch,  and  further,  along  with  nitrogen  and  other  elements,  into 
the  living  substance  of  plants,  and  finally  into  that  of  animals  as 
well,  since  the  latter  ultimately  depend  for  their  nourishment  ex- 
clusively on  plants.  As  a  result  of  the  progressive  simplification 
and  ultimate  oxidation  of  organic  materials,  man  and  other  animals 
obtain  the  energy  necessary  for  maintaining  the  body  temperature 
and  for  other  purposes. 

It  appears  therefore  that  the  carbon  in  nature  is  passing  through 
a  continuous  cycle,  being  alternately  reduced  in  the  green  parts 
of  plants  from  carbon  dioxide  to  compounds  relatively  rich  in 
carbon,  which  are  again  oxidized :  by  means  of  the  oxygen  of  the 
air  with  reformation  of  carbon  dioxide  and  water.  The  point  of 
fundamental  importance,  however,  is  the  energy  relations  of  these 
reactions.  Again  assuming  for  simplicity  that  the  compound  rich 
in  carbon  is  glucose,  the  thermochemical  equation  expressing  its 
oxidation  is  as  follows  : — 

C6H12O6+ 6O2->6CO2  +  6H2O  +  667,200  cal., 
1  This  oxidation  takes  place  both  in  plants  and  animals. 


352     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

that  is,  667,200  cal.  are  given  out  when  a  mol  of  glucose  is  com- 
pletely burned  to  carbon  dioxide  and  water.  The  reaction  therefore 
proceeds  spontaneously  in  the  direction  of  the  arrow  (p.  147),  and 
to  reverse  it  this  amount  of  energy  must  be  supplied.  This  energy 
comes  from  that  of  the  sun's  rays,  which,  by  a  method  not  yet  under- 
stood, is  absorbed  in  the  green  parts  of  plants,  and  stored  up  in  the 
form  of  compounds  relatively  rich  in  carbon  and  hydrogen. 

A  little  consideration  will  show  us  that  ultimately  we  depend 
almost  entirely  on  the  energy  thus  absorbed  by  the  green  parts 
of  plants  from  the  sun's  rays.  It  has  already  been  pointed  out 
that  our  nourishment  comes  directly  or  indirectly  from  plants, 
and  by  the  combustion  of  these  materials  animals  obtain  the 
energy  necessary  for  life.  Further,  the  energy  for  industrial  pro- 
cesses is  obtained  mainly  by  the  combustion  of  wood  and  of  coal. 
The  slow  change  of  wood  to  coal,  being  a  spontaneous  process, 
is  attended  by  a  loss  of  energy,  but  coal  still  represents  a  great 
amount  of  energy  in  a  compact  and  convenient  form.  It  is  not 
correct  to  say  that  the  energy  resides  in  the  coal  exclusively ; 
it  would  be  equally  correct  to  say  that  it  resides  in  the  oxygen,  as 
the  combination  of  the  two  is  essential  before  it  can  be  obtained. 
As,  however,  oxygen  is  everywhere  available,  the  only  expense 
incurred  is  in  obtaining  the  coal. 

The  nature  of  the  chemical  process  taking  place  in  the  green 
parts  of  plants  has  been  the  subject  of  numerous  investigations, 
but  has  not  been  definitely  elucidated.  One  view  is  that  carbon 
dioxide  and  water  react  to  produce  formaldehyde  and  oxygen  : 

C02+H20->CH20  +  02, 

the  former  then  undergoing  polymerization  to  form  a  simple  sugar  : 
6CH20->C6H12Oe. 

As  a  matter  of  fact,  traces  of  formaldehyde  can  be  detected  in  the 
green  parts  of  plants,  but  the  proof  that  the  first  steps  in  assimila- 
tion are  as  above  is  by  no  means  complete. 


CHAPTER  XXIII 
COMBUSTION   AND    FLAME 

/Combustion — Combustion  has  already  been  defined  as  the 
>~^  chemical  combination  of  two  substances  taking  place  with 
sufficient  vigour  to  develop  light  and  heat.  The  same  chemical 
action  may  or  may  not  take  place  with  combustion,  depending  on  the 
conditions.  Thus  hydrogen  and  oxygen,  in  contact  with  finely 
divided  platinum,  combine  slowly  without  sensible  rise  of  tempera- 
ture or  emission  of  light,  whilst  under  other  conditions  (p.  37)  the 
same  reaction  is  accompanied  by  a  flame  and  a  great  elevation  of 
temperature.  It  must  be  remembered,  however,  that,  in  accordance 
with  the  law  of  the  conservation  of  energy,  the  heat  given  out  in  a 
chemical  change  is  independent  of  the  way  in  which  the  change  is 
carried  out,  provided  that  the  final  products  are  the  same  in  each 
case. 

As  in  the  processes  of  combustion  with  which  we  are  most  familiar 
one  of  the  reacting  substances  is  the  oxygen  of  the  air,  the  substance 
capable  of  combining  vigorously  with  oxygen  is  said  to  be  combust- 
ible, and  the  atmosphere  is  said  to  be  a  supporter  of  combustion.  In 
the  same  way  other  gases,  such  as  nitrous  oxide  and  chlorine,  which 
behave  towards  combustible  substances  more  or  less  as  air  does,  are 
said  to  be  supporters  of  combustion.  Familiar  combustible  sub- 
stances are  coal,  coal  gas,  hydrogen,  sulphur,  hydrogen  sulphide,  etc. 
A  third  class  of  substances,  including  nitrogen,  carbon  dioxide,  and 
sulphur  dioxide,  are  neither  combustible  nor  supporters  of  combustion 
under  ordinary  conditions. 

A  little  consideration  shows  us,  however,  that  there  is  no  essential 
difference  between  a  combustible  substance  and  a  supporter  of  com- 
bustion. That  this  must  be  so  is  clear  when  it  is  realized  that  the 
essential  feature  of  the  process  is  chemical  combination  between  two 
substances,  and  both  are  equally  concerned  in  the  change.  This  is 
well  shown  by  the  arrangement  represented  in  Fig.  70,  in  which  air 
is  caused  to  burn  in  an  atmosphere  of  coal  gas.  An  ordinary  lamp- 
glass  is  fitted  at  the  lower  end  with  a  cork  carrying  a  central  glass 

23  353 


354     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


tube  and  two  side  tubes,  as  shown,  and  on  the  upper  opening  is  laid 
a  cover  with  a  round  hole  in  the  centre,  which  at  the  beginning  of  the 
experiment  is  closed  by  a  small  lid.  The  glass  is  filled  with  coal  gas 
by  the  side  tubes,  and  the  issuing  gas  ignited  at  the  lower  end  of  the 
central  tube.  The  lid  is  then  cautiously  removed  from  the  top  and 
the  issuing  gas  lighted.  The  draught  thus  caused  draws  the  flame 
to  the  upper  end  of  the  central  tube,  and  we  have  now  a  jet  of  air 
burning  in  coal  gas  in  the  interior,  and  coal 
gas  burning  in  air  at  the  top. 

The  interchangeability  of  the  terms  com- 
bustible and  supporter  of  combustion  is  further 
illustrated  by  the  fact  that  a  jet  of  hydrogen 
burns  in  chlorine  as  well  as  a  jet  of  chlorine 
in  hydrogen  (p.  90).  This  experiment  also 
illustrates  the  fact,  already  repeatedly  referred 
to,  that  the  term  combustion  is  not  confined 
to  reactions  in  which  oxygen  is  concerned. 

The  chemical  aspects  of  combustion,  and 
its  importance  for  the  development  of 
chemistry,  have  already  been  fully  discussed. 

Flame — When  the  combination  of  gaseous 
substances  takes  place  with  the  emission  of 
light  and  a  considerable  amount  of  heat,  the 
phenomenon  is  termed  a  flame.  The  restric- 
tion of  the  term  to  gaseous  substances  should 
be  noted.  When  carbon  combines  with 
oxygen  the  substance  glows  and  emits  light, 
but  there  is  no  flame.  The  appearance  of  a 
flame  when  certain  solid  substances  combine 
with  oxygen  is  due  to  the  intermediate  produc- 
tion of  gases  or  vapours.  In  the  case  of  sulphur 


FIG.  70. 


and  phosphorus,  for  instance,  the  heat  of  the  reaction  produces  vapours 
of  these  elements,  which  then  combine  with  oxygen,  giving  rise  to  a 
flame.  The  slight  bluish  flame  which  appears  over  burning  coal  is 
due  to  the  intermediate  formation  and  combustion  of  carbon  monoxide. 
The  same  explanation  applies  to  the  burning  of  a  candle,  combustible 
gases  being  continuously  produced  by  the  heat  of  combustion. 

Flames  differ  very  much  in  general  appearance.  The  flame  of 
hydrogen  or  of  alcohol  burning  in  air  is  practically  colourless,  that  of 
carbon  monoxide  is  blue,  and  that  of  cyanogen  lavender-coloured. 
Some  flames  have  a  very  high  degree  of  luminosity.  Thus  coal  gas 


COMBUSTION   AND   FLAME  355 

under  ordinary  conditions  burns  with  a  bright  yellowish  flame, 
and  acetylene  with  a  brilliant,  almost  white  flame.  The 
important  question  of  the  luminosity  of  flames  will  now  be 
considered. 

Luminosity  of  Flames — The  luminosity  of  the  flame  of 
magnesium  ribbon  burning  in  air  is  due  to  the  presence  of  particles 
of  magnesium  oxide,  which  become  incandescent  at  the  high  tempera- 
ture of  the  flame. 

The  luminosity  of  the  ordinary  coal-gas  flame  is  also  due  to  the 
presence  of  solid  particles — in  this  case  particles  of  carbon — heated 
to  incandescence.  The  evidence  often  adduced  in  favour  of  this 
view,  that  on  putting  a  cold  object — say  a  white  porcelain  dish — 
into  such  a  flame  particles  of  soot  are  deposited,  is  not  very  con- 
vincing. More  conclusive  is  the  fact  that  if  a  coal-gas  flame 
is  placed  between  a  bright  light  and  a  screen  the  luminous 
portion  of  the  flame  casts  a  shadow  on  the  screen,  and,  further,  if 
the  image  of  the  sun  is  focused  upon  the  luminous  part  of  a  flame 
the  scattered  light  is  found  to  be  polarized. 

It  has  already  been  stated  that  the  luminosity  of  the  coal-gas  flame 
is  connected  with  the  presence  in  coal  gas  of  unsaturated  hydrocarbons, 
such  as  ethylene  and  acetylene.  We  now  understand  that  in  the 
process  of  combustion  these  gases  must  at  some  stage  give  rise  to 
free  carbon,  upon  the  presence  *of  which  the  luminosity  depends. 
Methane,  on  the  other  hand,  burns  without  the  intermediate  pro- 
duction of  free  carbon.  Coal  gas  of  small  luminosity  can  be  prepared 
more  cheaply  than  a  gas  of  high  luminosity,  and  as  the  former  is 
equally  efficient  for  heating  purposes,  it  is  now  largely  used  in 
commerce,  being  "enriched"  by  the  addition  of  unsaturated  hydro- 
carbons when  required  for  illuminating  purposes.  The  tendency  to 
the  commercial  production  of  a  cheaper,  slightly  luminous  gas  has 
received  further  support  by  the  extended  use  of  incandescent  mantles 
(see  below). 

The  presence  of  solid  particles  is  not,  however,  the  sole  cause  of 
luminosity  in  flames.  Thus  the  vapour  of  carbon  disulphide  burning 
in  oxygen  or  in  nitric  oxide  gives  a  brilliant  flame,  although  no  solid 
matter  is  present.  The  luminosity  in  such  cases  is  ascribed  by  some 
investigators  to  "  luminescence."  1 

The  luminosity  of  flames  may  often  be  increased  (i)  by  increasing 

1  The  emission  of  light  arising  from  the  temperature  possessed  by  a  body  is 
termed  incandescence;  emission  of  light  arising  from  all  other  causes  than  tempera- 
ture is  termed  luminescence. 


356     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


the  pressure;  (2)  by  increasing  the  temperature.  Thus  Frankland 
showed  that  the  flame  of  hydrogen  burning  in  oxygen,  which  is 
practically  non-luminous,  becomes  luminous  when  the  pressure  is 
increased.  The  effect  of  temperature  on  the  luminosity  is  taken 
advantage  of  in  the  so-called  recuperative  burners,  in  which  the  gases 
are  warmed  before  being  introduced  into  the  flame. 

It  is  a  familiar  fact  that  when  lime  is  raised  to  a  high  temperature 
it  glows  and  emits  a  very  bright  light.  The  same  principle  has  been 
largely  used  in  recent  years  in  the  construction  of  incandescent 
mantles,  which  consist  simply  of  a  skeleton  of  infusible  oxides  heated 
to  incandescence  in  a  non-luminous  flame  (Bunsen  flame).  For 
reasons  as  yet  unexplained,  a  mixture  of  oxides  gives  much  better 
results  than  any  one  oxide.  The  mixture  most  largely  used  consists 
of  99  per  cent,  thorium  dioxide  and  I  per  cent,  cerium  dioxide. 

The  luminosity  of  the  ordinary  gas  flame  is  diminished  (i)  by 
cooling  it,  for  example,  by  previously  adding  carbon  dioxide  or 
nitrogen  to  the  gas;  (2)  by  preventing  the  liberation  of  carbon 
particles  by  previously  mixing  the  coal  gas  with  oxygen.  The  latter 
method  is  used  in  the  Bunsen  flame  (g.v.}. 

Structure  of  Flame.  The  Bunsen  Flame— The  ordinary 
coal-gas  flame  (or  candle  flame)  consists  of  four  distinct  regions 
(Fig.  71). 

(i)  The  inner  dark  zone  a,  consisting  of  unburnt  and  practically 


--a 


FIG.  71. 


FIG.  72. 


unaltered  gas.  This  may  be  proved  by  abstracting  some  of  the  gas 
by  means  of  a  glass  tube  inserted  in  the  flame  and  lighting  it 
(Fig.  72) 


COMBUSTION   AND   FLAME 


357 


(2)  Surrounding  the  inner  zone  is  the  luminous  region  £,  in  which 
partial    combination   occurs.     The  un- 

saturated  hydrocarbons  (and  perhaps 
other  hydrocarbons  as  well)  are  decom- 
posed at  the  relatively  high  temperature 
of  this  region  with  liberation  of  carbon 
particles,  the  glowing  of  which  is  the 
source  of  the  luminosity  of  the  flame. 

(3)  In  the  outer  slightly  bluish  margin 
<r,  where  excess  of  oxygen  is  available, 
complete  combustion  to  carbon  dioxide 
occurs. 

(4)  The  bluish  region,  d,  at  the  base 
of  the   flame.      In   this   region,   which 
appears  to  correspond  with  the  bluish- 
green  sheath  of  the  Bunsen  flame,  but 

does  not  extend  over  the  whole  flame,  FlG-  73- 

the  partial  combination  may  reach  the  same  stage  as  in  the  region  of 
the  Bunsen  flame  just  mentioned  (see  below). 

The  principle  of  the  Bunsen  burner,  already 
referred  to,  is  illustrated  in  Fig.  73.  The  gas  issues 
from  a  small  jet  d  at  the  bottom  of  the  metal  tube  b, 
and  in  passing  upwards  causes  a  reduction  of  pres- 
sure by  means  of  which  air  is  drawn  in  at  the 
openings  on  either  side  of  the  jet.  The  air  and  gas 
mix  on  their  way  up  the  tube  and  burn  at  the  top 
with  a  practically  non-luminous  (slightly  bluish) 
flame.  Under  ordinary  circumstances  the  amount 
of  air  taken  in  is  considerably  less  than  that  re- 
quired for  complete  combustion  of  the  coal  gas. 
The  flame  maintains  its  position  when  the  velocity 
with  which  the  mixture  of  gas  and  air  issues  from 
the  tube  exceeds  its  velocity  of  inflammation  ;  if  the 
velocity  is  less  the  flame  "  strikes  back "  and  con- 
tinues to  burn  at  the  base  of  the  tube. 

The  Bunsen  flame  (Fig.  74)  differs  mainly  from 
the  flame  described  above  in  that  the  luminous 
region  is  absent.  It  consists  of  three  regions  (i)  the 
inner  cone  «,  of  unburnt  gas  and  air  ;  (2)  a  narrow, 
bluish-green  sheath  <:,  surrounding  the  inner  cone,  in  which  the  gas 
and  air  drawn  in  react  to  form  a  mixture  of  nitrogen  (from  the  air), 


FIG.  74. 


358     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


carbon  monoxide,  carbon  dioxide,  water  vapour,  and  hydrogen ; 
(3)  the  practically  non-luminous  outer  mantle  £,  where  the  gases 
from  the  inner  cone  finally  burn  completely  to  carbon  dioxide  and 
water. 

In  the  intermediate  region  there  is  practically  no  oxygen  and  a 
fairly  high  temperature,  so  that  substances  held  in  this  region  are 
readily  reduced.  In  the  outer  mantle,  on  the  contrary,  there  is  excess 
of  hot  oxygen,  and  it  therefore  acts  as  an  oxidizing  flame. 

Investigation  of  Chemical  Changes  taking  place  in 
Flames — It  is  evidently  not  quite  an  easy  matter  to  investigate 
experimentally  the  changes  taking  place  in  the  different  regions  of 
the  Bunsen  or  any  other  flame.  An  arrangement  for  withdrawing 

the  gases  from  different  levels  of  the 
flame  may  itself  cause  disturbances  and 
modify  the  changes.  This  difficulty  is 
to  some  extent  got  over  by  means  of  an 
arrangement  introduced  independently 
by  Teclu  and  by  Smithells,  known  as 
aflame-separator.  It  consists  of  a  wide 
glass  tube  which  is  fitted  tightly  over 
the  upper  part  of  a  narrow  tube,  as 
shown  in  Fig.  75.  The  gas  and  air  are 
admitted  separately  at  the  base  of  the 
inner  tube,  and  by  suitably  regulating 
the  supply  of  each  the  external  bluish 
cone  is  made  to  burn  at  the  upper  end 
of  the  wide  tube  while  the  green  inner 
cone  burns  at  the  top  of  the  inner  tube. 
By  means  of  a  side  tube,  #,  on  the  wide 
glass  tube,  the  interconal  gases  can  be 
withdrawn  and  analyzed,  and  in  this 
way  the  nature  of  the  chemical  action 
in  the  greenish  cone  has  been  definitely 
established. 

Owing  to  the  experimental  difficulties 
of  the  investigation,  the  exact  changes 
taking  place  in  luminous  flames,  more 
particularly  the  appearance  of  free  carbon,  are  not  well  understood, 
t  was  once  thought  that  the  liberation  of  carbon  was  due  to  the 
preferential  combustion  of  hydrogen,  but  this  view  is  now  abandoned. 
A  more  plausible  explanation  is,  that  at  the  high  temperature  of  the 


FIG.  7S. 


COMBUSTION   AND   FLAME  359 

flame  ethylene  dissociates  into  acetylene  and  hydrogen  and  the 
former  into  free  carbon  and  hydrogen  : 

C2H4->C2H2+  H2 ;  C2H2->2C  +  H2  ; 

in  other  words,  dissociation  precedes  combustion.  A  modification  of 
this  view  is  that  hydrocarbons  yield  free  carbon  by  thermal  decom- 
position without  the  intermediate  production  of  acetylene.  Bone,  on 
the  other  hand,  is  of  the  opinion  that  the  hydrocarbons  first  unite  with 
oxygen  forming  "  hydroxylated  "  compounds  which,  in  the  absence  of 
excess  of  oxygen,  break  down  into  free  carbon  and  other  products.  It 
would  lead  too  far  to  discuss  this  interesting  subject  more  fully  here. 

Temperature  Of  Flames — The  temperature  of  any  particular 
flame  is  not  quite  constant,  as  it  depends  on  the  condition  of  the 
combustible  substance,  on  the  amount  of  oxygen  supplied,  etc.  The 
temperature  of  the  green  inner  cone  of  the  Bunsen  flame  coal-gas-air 
is  about  1500°,  that  of  the  outer  cone  about  1800°.  The  temperature 
of  the  outer  cone  when  oxygen  is  supplied  instead  of  air  is  2200°. 
The  ordinary  acetylene  flame  gives  a  temperature  of  about  1900°, 
when  a  Bunsen  burner  is  used- it  rises  to  2500°.  The  hydrogen- 
oxygen  flame  gives  a  temperature  of  about  2400°,  that  of  carbon 
monoxide  burning  in  oxygen 
as  high  as  2800°.  The  num- 
bers refer  to  the  maximum 
temperatures  observed. 

It  might  be  anticipated 
that  the  temperature  of  a  ' 
flame  could  be  calculated 
from  the  heat  of  combustion  of  the  products  and  their  respective  heat 
capacities  (specific  heats),  provided  that  due  allowance  is  made  for  the 
loss  of  heat  by  radiation.  Tailing  as  a  simple  illustration  the  hydrogen- 
oxygen  flame,  the  heat  of  formation  of  i  mol  (18  grams)  of  water  is 
about  68,400  calories,  and  as  the  specific  heat  of  water  vapour  is  about 
9  calories  per  mol,  we  obtain  for  the  calculated  temperature  68,400/9 
=  7,600°.  Even  if  allowance  is  made  for  the  heat  lost  by  radiation, 
this  enormously  exceeds  the  observed  value.  The  explanation  is  that 
above  a  certain  temperature  the  combination  of  hydrogen  and  oxgen 
is  no  longer  complete,  and  the  higher  the  temperature  the  less  complete 
the  combination  (p.  52).  The  further  combustion  is  only  rendered 
possible  by  the  loss  of  heat  owing  to  radiation,  and  the  observed 
temperature  is  a  kind  of  equilibrium  between  these  different  factors. 

A  suitable  arrangement  for  obtaining  a  flame  with  two  gases,  e.g. 
hydrogen  and  oxygen,  is  illustrated  in  Fig.  76.     One  gas  passes  along 


360    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

a  moderately  wide  metal  tube,  the  other  is  conveyed  by  a  narrow 
metal  tube  in  the  axis  of  the  wider  tube,  and  the  tubes  are  so  arranged 
that  the  gases  mix  only  just  before  reaching  the  flame. 

Ignition  Temperature.  The  Davy  Lamp— The  rate  of 
combination  of  a  mixture  of  hydrogen  and  oxygen,  like  that  of  other 
chemical  reactions,  depends  greatly  on  the  temperature.  Below  a 
certain  temperature  the  change  is  comparatively  slow,  but  above  this 
temperature  it  takes  place  with  explosion.  Similarly  phosphorus  is 
slowly  oxidized  in  the  air  at  room  temperature,  but  when  the  tem- 
perature is  sufficiently  high,  the  combination  is  so  vigorous  that  heat 
and  light  are  given  out ;  in  other  words  combustion  occurs.  The 
lowest  temperature  which  enables  explosion  or  combustion  to  take  place 
in  a  system  is  termed  the  ignition  temperature. 

The  occurrence  of  an  ignition  point  in  a  system  is  closely  connected 
with  the  heat  given  out  in  a  reaction.  Above  the  ignition  point, 
combination  of  hydrogen  and  oxygen  is  so  rapid  that  a  large  amount 
of  heat  is  given  out  in  a  short  time  ;  this  raises  a  further  quantity 
of  the  gas  above  the  ignition  temperature,  this  combines  with  emission 
of  more  heat,  and  so  on,  combination  spreading  quickly  through  the 
whole  system.  Below  the  ignition  point,  however,  combination  is 
slow  ;  the  heat  given  out  is  dissipated  over  a  wide  area  and  the 
temperature  never  reaches  that  required  for  ignition.  It  does  not 
follow,  however,  that  explosion  or  combustion  will  necessarily  occur 
when  the  temperature  is  raised  at  one  point  above  that  required  for 
rapid  combination  ;  this  again  depends  upon  the  thermal  behaviour 
of  the  system.  If  the  heat  given  out  in  combustion  is  sufficient  to 
raise  the  system  above  the  ignition  temperature,  combustion  will 
proceed  once  it  has  started  without  the  further  application  of  heat. 
This  can  of  course  only  occur  in  exothermic  reactions,  and  numerous 
examples  have  already  been  given.  On  the  other  hand,  if  the  heat 
of  combination  is  less  than  that  required  to  raise  the  system  above 
the  temperature  of  ignition,  combination  cannot  proceed  unless  energy 
is  continuously  supplied  from  without.  The  combination  of  nitrogen 
and  oxygen  under  the  influence  of  the  electric  discharge  (p.  223)  is 
an  illustration  of  this  case.  On  these  considerations  is  based  an 
alternative  definition  of  the  ignition  temperature  as  the  temperature 
at  which  the  initial  loss  of  heat  due  to  conduction^  etc.^  is  equal  to  the 
heat  evolved  in  the  same  time  by  the  chemical  reaction  (van  't  Hoflf). 

The  experimental  difficulty  of  determining  ignition  temperatures 
with  accuracy  is  very  considerable.  When  solids  are  concerned,  the 
state  of  division  is  of  importance,  and  in  the  case  of  gaseous  mixtures 


COMBUSTION   AND   FLAME 


36: 


the  results  are  influenced  by  the  positive  or  negative  action  of  the 
walls  of  the  containing  vessel.  Thus  the  ignition  temperature  of 
phosphorus  in  air  is  usually  given  at  35°,  but  we  have  seen  that  when 
deposited  in  a  finely  divided  form  from  solution  in  carbon  disulphide, 
it  catches  fire  spontaneously  at  room  temperature.  The  uncertainty 
of  the  results  obtained  with  gases  is  well  illustrated  by  the  fact  that 
the  values  given  in  the  literature  for  the  ignition  temperature  of 
electrolytic  gas  (2H2 :  O2)  vary  from  518°  to  845°. 

The  difficulties  in  the  case  of  gaseous  mixtures  have  been  largely 
overcome  by  an  ingenious  method  due  to  Dixon,  who  has  succeeded 
in  igniting  the  gases  out  of  contact  with  any  solid  material.  One 
gas  was  passed  through  a  wide  porcelain  tube,  the  other  through  a 
narrow  tube  fixed  along  the  axis  of  the  wide  tube.  The  temperature 
was  gradually  raised  till  ignition  occurred  ;  this  took  place  at  a 
point  above  the  orifice,  and  special  experiments  showed,  as  was  to 
be  anticipated,  that  the  ignition  temperatures  thus  determined  were 
independent  of  the  materials  of  the  tubes. 

The  mean  ignition  temperatures  for  a  few  of  the  commoner  gases 
at  normal  pressure  are  as  follows  : — 


In  Oxygen. 

In  Air. 

Hydrogen  .... 

5850 

585° 

Carbon  monoxide  (moist)  . 

650° 

65I° 

Acetylene  .... 

428° 

429° 

Hydrogen  sulphide    . 

227° 

364° 

As  the  table  shows,  the  ignition  temperatures  of  the  first  three  gases 
are  the  same  in  air  and  in  oxygen  within  the  limits  of  experimental 
error  ;  that  of  hydrogen  sulphide  is  much  ^ 

higher  in  air. 

The  ignition  point  of  a  substance  in  air      ^ 
may  be  below  room  temperature  ;  it  is  then    w;* 
said  to  be  .spontaneously  inflammable.     Ex- 
amples of  this  have  already  been  given. 

If  a  piece  of  wire  gauze  be  held  a  little 
above  a  Bunsen  burner  and  the  gas  lighted 
on  the  upper  side  (Fig.  77),  the  flame  does 
not  pass  through  to  the  gas  on  the  lower  side.     This  phenomenon  is 
connected  with  the  fact  that  wire  gauze  is  an  excellent  conductor  of 


FIG.  77. 


362     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

heat,  and  the  explanation  usually  given  is  that  so  much  of  the  heat 
is  conveyed  away  by  the  gauze,  that  the 
mixture  of  gas  and  air  on  the  lower  side  of 
the  gauze  does  not  reach  the  ignition  tem- 
perature. 

On  this  principle  is  based  the  well-known 
Davy  safety-lamp  used  by  miners.  It  con- 
sists of  an  ordinary  oil  lamp,  the  flame  of 
which  is  enclosed  in  a  glas=>  cylinder  sur- 
mounted by  a  cylinder  of  wire-gauze  (Fig.  78). 
When  this  lamp  is  used  in  a  mine  the 
atmosphere  of  which  contains  methane,  the 
flame  does  not  ignite  the  explosive  mixture 
outside,  for  the  reason  given  above.  The 
methane,  however,  passes  through  the  gauze 
and  burns  inside,  causing  considerable 
alteration  in  the  appearance  of  the  flame. 
From  the  amount  of  alteration  thus  pro- 
duced, the  proportion  of  methane  in  the 
atmosphere  of  the  mine  may  be  roughly 
estimated.  If  part  of  the  flame  is  acci- 
dentally blown  through  the  gauze,  or  even 
FlG-  ?8-  if  the  latter  becomes  strongly  heated  at  any 

point,  explosion  may  occur. 

Calorific  Value  of  Fuels — The  calorific  value  of  a  fuel  (cf.  p.  349)  may  be 
defined  as  the  amount  of  heat  obtained  by  the  combustion  of  unit  weight  of  the 
fuel.  The  estimation  is  most  conveniently  made  by  burning  a  known  weight  of 
the  material  in  compressed  oxygen  in  some  form  of  bomb  calorimeter,  the  Berthelot 
apparatus1  being  very  suitable.  This  method  gives  the  heat  of  combustion  at 
constant  volume,  which  can  be  corrected  to  constant  pressure  if  desired  (p.  146). 

As  the  heat  of  combustion  of  amorphous  carbon  is  96,980  calories,  the  calorific 
value  of  charcoal,  expressed  according  to  the  above  definition,  is  8080  calories. 
The  calorific  value  of  some  important  constituents  of  fuels  is  given  in  the  accom- 
panying table  (cf.  p.  338). 


Calories  per  gram. 
Carbon,  amorphous      .     8,080 
Anthracite  coal      .        .     9,200-9,800 
Carbon  to  CO       .        .    2,410 
Carbon  monoxide          .    2,430 


Calories  per  gram. 

Hydrogen  (to  liquid  water)    .     34,200 
Methane  „  .     13,250 

Ethylene  , ,  .     11,900 

Acetylene  ,,  .     11,900 


The  calorific  value  of  a  fuel  may  be  expressed  in  terms  of  calories  per  gram- 
molecule  or  in  other  ways,  and  that  of  a  gaseous  fuel  may  also  be  expressed  in 
terms  of  calories  per  unit  volume. 

1  Physical  Chemistry,  p.  147. 


CHAPTER   XXIV 
SILICON   AND    BORON 

WITH  the  exception  of  silicon  and  boron,  which  are  discussed 
in  the  present  chapter,  all  the  typical  non-metals  have  now 
been  considered.      Silicon  and   boron  do  not   belong  to  the  same 
group  of  elements,  but  some  of  their  compounds  show  considerable 
analogy. 

SILICON 

Symbol,  Si.     Atomic  weight=28.3.     Molecular  weight  unknown. 

• 
Silicon  belongs  to  the  carbon  group  of  elements,  and  appears  to 

function  invariably  as  a  quadrivalent  element  As  will  appear  later, 
there  is  a  close  analogy  between  carbon  and  silicon,  and  this  applies 
to  the  elements  themselves  as  well  as  to  their  compounds. 

Occurrence  —  Silicon  does  not  occur  free  in  nature.  In  com- 
bination with  oxygen,  as  silicon  dioxide,  SiO2,  and  with  oxygen  and 
different  metals,  as  silicates,  it  is  very  widely  distributed.  Quartz, 
flint,  and  white  sand  are  nearly  pure  silicon  dioxide.  Silicates  enter 
largely  into  the  composition  of  rocks  and  the  different  varieties  of 
clay,  and  are  present  in  the  soil,  from  which  they  are  taken  up  by 
plants.  Next  to  oxygen,  silicon  is  the  most  abundant  element  in  the 
earth's  crust. 

Preparation  —  Silicon,  like  carbon,  occurs  in  different  allotropic 
modifications,  an  amorphous  and  at  least  one  crystalline  form  being 
known. 

(i)  Amorphous  silicon  is  obtained  by  heating  potassium  silico- 
fluoride  (p.  366)  with  metallic  potassium  : 


(2)  The  same  modification  is  obtained  in  still  purer  condition  by 
heating  silicon  dioxide  (powdered  quartz,  sand,  or  better,  the  pure 
precipitated  oxide)  with  metallic  magnesium  : 

SiO2  +  2  Mg->2  MgO  +  Si. 
363 


364     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

If  excess  of  magnesium  is  used,  as  is  advantageous,  some  magnesium 
silicide,  Mg2Si,  is  formed,  but  it  and  the  oxide  are  removed  by  treat- 
ment with  hydrochloric  acid,  and  fairly  pure  silicon  remains. 

(3)  The  same  form  is  obtained  by  passing  the  vapour  of  silicon 
tetrachloride  over  heated  metallic  sodium  : 


The  sodium  chloride  is  removed  by  washing  with  water. 

(4)  Silicon  is  obtained   in    octahedral   crystals   by   heating   in  a 
crucible  3  parts  of  potassium  silicofluoride,  I  part  of  sodium,  and  i 
part  of  zinc,  and  dissolving  out  the  other  products  by  treatment  with 
acid.      This  method  of  preparation  depends  upon  the  fact  that  the 
silicon  set  free  by  the  action  of  metallic  sodium  on  the  silicofluoride 
dissolves  in  fused  zinc,  and  separates  in  crystalline  form  on  cooling. 

(5)  Crystalline  silicon  is  now  prepared  on  the  commercial  scale  by 
heating  quartz  sand  with  coke  in  the  electric  furnace  : 


Properties  of  Amorphous  Silicon  —  Amorphous  silicon  is 
a  brown  powder  of  density  2.35.  It  burns  when  heated  with  air  or 
oxygen  to*  the  dioxide,  but  the  latter  compound,  being  non-volatile, 
coats  the  particles  of  the  element  and  retards  the  action.  It  com- 
bines directly  with  fluorine  at  the  ordinary  temperature,  with  chlorine 
at  450°,  with  bromine  at  450°,  and  with  iodine  at  a  red  heat,  forming 
compounds  of  the  type  SiX4.  It  combines  directly  with  carbon 
(p.  367),  with  boron,  and  other  elements  when  heated  in  the  electric 
furnace.  It  is  a  powerful  reducing  agent  at  high  temperatures  ;  for 
example,  it  reduces  phosphorus  pentoxide  to  phosphorus.  It  is 
readily  dissolved  by  liquid  hydrogen  fluoride,  with  formation  of 
hydrofluosilicic  acid,  H2SiF6,  and  evolution  of  hydrogen: 

Si  +  6HF->SiF4,2HF  +  2H2. 

The  aqueous  solutions  of  acids,  even  hydrofluoric,  have  very  little 
action,  but  it  is  readily  dissolved  by  a  mixture  of  nitric  and  hydro- 
fluoric acids.  It  is  readily  dissolved  on  boiling  with  sodium  or 
potassium  hydroxide,  alkali  silicate  and  hydrogen  being  formed  : 

Si+2NaOH  +  H2O->Na2Si03+2H2. 

Properties  of  Crystalline  Silicon  —  This  modification 
occurs  in  dark  grey,  opaque  octahedral  crystals  somewhat  resembling 
graphite  ;  the  density  varies  from  2.  i  to  2.49.  Crystalline  silicon  con- 


SILICON   AND    BORON  365 

ducts  electricity  slightly;  the  amorphous  form  has  no  conducting 
power.  The  crystals  are  very  hard,  readily  scratching  glass. 

In  chemical  behaviour,  the  crystalline  resembles  the  amorphous 
form,  but  being  in  a  less  finely  divided  condition  is  less  active.  The 
crystalline  form  is  considerably  more  stable  towards  oxygen  than  the 
other  modification. 

Hydrogen  Silicides  —  Two  compounds  of  silicon  and  hydrogen, 
SiH4  and  Si2H6,  the  analogues  of  methane  and  ethane  respectively,  are 
known. 

Hydrogen  Silicide  or  Siiicoinethane,  SiH4—  Prepara- 
tion —  When  magnesium  silfcide,  Mg2Si,  prepared  as  already 
described  (p.  364),  is  treated  with  concentrated  hydrochloric  acid, 
a  spontaneously  inflammable  mixture  containing  the  two  hydrogen 
silicides  and  free  hydrogen  is  obtained.  In  order  to  obtain  the  pure 
compounds,  the  gaseous  mixture  is  passed  through  a  tube  dipping  in 
liquid  air,  when  the  two  silicides  are  condensed  and  the  hydrogen 
passes  on.  The  silicides  are  then  separated  by  fractional  distillation 
(cf.  p.  208). 

Properties  —  Silicomethane  is  a  colourless  gas,  which  is  not 
spontaneously  inflammable  at  atmospheric  pressure,  but  i»flames 
when  the  pressure  is  reduced  or  when  it  is  slightly  warmed  ;  it  then 
burns  with  a  bright  flame  to  silicon  dioxide  and  water.  When  passed 
into  the  solution  of  an  alkali,  there  is  a  vigorous  evolution  of  hydrogen, 
and  an  alkali  silicate  is  formed  : 

SiH4+  2KOH  +  H2O-»K2SiO3+  4H2. 

When  heated  above  400°,  it  decomposes  into  its  elements. 

Silicoethane,  Si2H6  —  The  preparation  of  this  compound  has 
been  described  above. 

Properties  —  Silicoethane  is  a  colourless  liquid,  heavier  than 
water  ;  it  boils  at  52°,  and  the  crystals  melt  at  -  138°.  It  is  spontane- 
ously inflammable  in  air  at  the  ordinary  temperature,  burning  with  a 
brilliant  flame,  silicon  dioxide  and  amorphous  silicon  being  deposited. 
It  is  completely  decomposed  into  its  elements  on  heating/to  250°.  On 
passing  into  sodium  hydroxide  solution,  hydrogen  is  given  off  and 
sodium  silicate  formed  : 


COMPOUNDS  OF  SILICON  WITH  THE  HALOGENS 

Silicon  Fluoride,  $\¥  ^Preparation—  (i)  By  direct  combina- 
tion of  silicon  and  fluorine  (Moissan). 


366     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

(2)  By  heating  silicon  dioxide  (sand)  with  hydrofluoric  acid,  or, 
more  conveniently,  by  heating  together  calcium  fluoride,  sand  and 
sulphuric  acid.  In  the  latter  case,  hydrofluoric  acid  is  obtained 
on  heating,  and  immediately  attacks  the  silicon  dioxide  : 

+  H2SO4->CaSO4+2HF 


The  etching  effect  of  hydrofluoric  acid  on  glass  (p.  153)  depends  on 
this  reaction. 

(3)  Pure  silicon  fluoride  is  obtained  by  heating  dry  barium  silico- 
fluoride  : 

BaSiF6->BaF2+  SiF4. 

The  gas  can  be  collected  over  mercury  in  the  entire  absence  of 
moisture. 

Properties  —  Silicon  fluoride  is  a  colourless  gas,  with  a  suffocating 
odour  ;  it  fumes  in  contact  with  moist  air.  It  is  stable  even  at 
high  temperatures,  corresponding  with  its  great  heat  of  formation, 
23,980  cal.,  and  is  not  decomposed  even  by  electric  sparks.  When 
passed  into  water  it  is  decomposed,  with  formation  of  hydrofluosilicic 
acid,  which  remains  in  solution,  and  silicic  acid,  which  is  precipi- 
tated : 


As  the  insoluble  silicic  acid  would  soon  stop  up  the  tube  conveying 
the  fluoride,  the  experiment  is  best  performed  by  pushing  the  end 
under  mercury.  As  each  bubble  of  gas  escapes  and  comes  into  contact 
with  water,  it  is  decomposed  with  precipitation  of  silicic  acid. 

Hydrofluosilicic  Acid,  H2SiF6—  When  the  precipitated  silicic 
acid  is  removed  by  filtration,  an  aqueous  solution  of  hydrofluosilicic 
acid,  H2SiF6,  is  obtained.  The  acid  is  not  known  in  the  pure  condi- 
tion ;  when  a  concentrated  aqueous  solution  is  evaporated,  silicon 
fluoride  escapes  most  rapidly  and  hydrofluoric  acid  remains.  The 
salts  are  decomposed  in  an  analogous  way  on  heating  (see  (3)  above). 
Barium  silicofluoride  is  practically  insoluble  in  water,  and  this  pro- 
perty is  taken  advantage  of  as  a  test  both  for  barium  and  for  fluosili- 
cates.  Potassium  fluosilicate  is  only  slightly  soluble  in  water  (i  in  833 
at  17.5°).  Fluosilicates  are  decomposed  in  a  rather  remarkable  way 
by  alkalis  in  the  cold,  alkali  fluoride  being  formed  and  silicic  acid 
precipitated  : 

aSiOs+  H2O. 


SILICON   AND   BORON  367 

Silicon  Tetrachloride,  S\C\—Preparatioji~(i)  By  heating 
either  amorphous  or  crystalline  silicon  in  a  stream  of  chlorine. 

(2)  By  the  action  of  chlorine  on  a  heated  mixture  of  silicon 
dioxide  and  carbon: 


Neither  carbon  nor  chlorine  alone  have  any  effect  on  silicon 
dioxide  under  the  conditions  of  the  above  experiment.  The  prin- 
ciple of  the  method  is  that  by  using  both  substances  much  more 
energy  is  given  out  than  would  be  the  case  if  silicon  or  oxygen  were 
products  of  the  reaction  (p.  147).  The  method  is  also  used  for  other 
chlorides. 

Properties  —  Silicon  tetrachloride  is  a  colourless  liquid  with  a 
suffocating  odour;  it  boils  at  57°,  and  its  density  at  o°  is  1.52.  It 
fumes  in  moist  air,  and  is  completely  decomposed  by  water  : 

SiCl4  +  3H2O-»4HCl+H2SiO3. 

When  silicon  is  heated  in  a  stream  of  hydrogen  chloride,  silicon 
chloroform,  SiHCl3,  the  silicon  analogue  of  chloroform,  is  obtained. 
It  is  a  colourless  liquid  which  boils  at  33°. 

Silicon  Tetrabromide,  SiBr4,  and  the  tetraiodide,  SiI4,  can 
be  obtained  by  methods  analogous  to  those  described  in  connexion 
with  the  chloride.  The  former  is  a  colourless  liquid  which  boils  at 
130°,  the  iodide  occurs  in  colourless,  octahedral  crystals,  melting 
at  120.5°. 

Carborundum,  CSi,  is  prepared,  according  to  Acheson,  by  heat- 
ing together  in  the  electric  furnace  a  mixture  of  silicon  dioxide  (sand), 
coke  and  common  salt. 

Properties  —  Carborundum  occurs  in  hexagonal  crystals,  which  are 
colourless  when  pure,  the  density  is  3.2.  It  is  not  affected  by  any 
acid  or  mixture  of  acids,  is  scarcely  affected  by  oxygen  even  at  1000°, 
but  is  decomposed  by  fused  alkalis.  Next  to  the  diamond  and  boron 
carbide  it  is  the  hardest  substance  known,  and  is  largely  used  for 
grinding  and  polishing  purposes. 

Silicon  Dioxide,  SiO2—  Occurrence—  As  already  indicated, 
silicon  dioxide  is  very  widely  distributed  in  nature  ;  in  fact  the  greater 
part  of  the  crust  of  the  earth  is  composed  of  this  compound,  and  the 
salts  of  the  corresponding  acids,  the  silicates. 

Silicon  dioxide  occurs  in  two  chief  crystalline  modifications  —  (a) 
quartz  with  its  numerous  varieties,  including  rock  crystal,  amethyst, 
smoky  quartz,  etc.  ;  (&}  tridymite  ;  and  also  in  the  amorphous  form  as 


368     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

flint,  opal,  jasper,  kieselguhr,  etc.  It  is  also  found  in  most  plants,  and 
is  particularly  abundant  in  the  stems  of  grasses,  equisetums  (horse- 
tails), bamboo,  etc. 

Modes  of  Formation — (i)  It  is  obtained  by  burning  silicon  in  air. 

(2)  Pure  amorphous  silicon  dioxide  is  obtained  by  igniting  silicic 

acid  (g.v.) : 

H2SiO3->H2O4-SiO2. 

Properties — The  purest  form  of  quartz,  rock  crystal,  occurs  in  per- 
fectly colourless,  six-sided  crystals  of  density  2.6.  Some  of  the 
crystals  rotate  polarized  light  to  the  right,  others  to  the  left,  and  this 
property  is  taken  advantage  of  in  the  construction  of  polarimeters. 
Common  quartz  is  white,  but  not  transparent,"  the  varieties  already 
mentioned,  including  amethyst,  smoky  quartz,  sand-stone,  etc.,  owing 
their  colours  to  traces  of  impurities.  Owing  to  its  great  hardness, 
quartz  (especially  in  the  form  of  sandstone)  is  used  as  a  grinding 
material.  Quartz  is  only  stable  up  to  900° ;  above  this  temperature  it 
changes  to  tridymite.  The  latter  is  found  in  small  crystals  in  certain 
rocks,  its  density  is  only  2.3.  It  can  be  obtained  both  in  hexagonal 
and  in  rhombic  crystals. 

The  amorphous  forms  of  silicon  dioxide,  including  flint,  opal,  agate 
and  jasper,  owe  their  colours  to  traces  of  impurities,  often  oxide  of 
iron.  Kieselguhr  and  the  silicon  dioxide  obtained  by  ignition  are 
practically  colourless. 

Crystalline  and  amorphous  silicon  dioxide  begin  to  soften  when 
heated  to  1500°,  and  are  completely  liquid  at  1750°.  Pieces  of  chemical 
apparatus  are  now  obtainable  prepared  by  fusing  quartz  in  the  oxy- 
hydrogen  flame.  Quartz  is  remarkable  for  its  extremely  low  co- 
efficient of  expansion,  and  therefore  a  vessel  made  of  silica  may  be 
made  red-hot  and  dropped  into  cold  water  without  risk  of  fracture. 

Silicon  dioxide  is  not  affected  by  water,  or  by  acids  other  than 
hydrofluoric  acid  (p.  153).  When  fused  with  alkalis  or  alkali  car- 
bonates, silicates  are  formed : 

SiO2  +  2Na2CO3->Na4SiO4  +  2CO2. 

Silicic  Acids — As  silicon  dioxide  is  an  acidic  oxide,  the  exist- 
ence of  silicic  acids  of  the  types,  H2O.SiO2  or  H2SiO3  ;  SiO2,2H2O, 
or  H4SiO4,  and  so  on,  might  be  anticipated.  As  a  matter  of  fact, 
there  appears  to  be  a  whole  series  of  such  acids  of  the  general 
formula,  ;rSiO2vyH2O,  where  x  andjj/  are  whole  numbers.  Up  to  the 
present  none  of  them  has  been  obtained  in  a  pure  condition  ;  but  salts 
derived  from  them,  the  silicates,  form  the  chief  constituents  of  rocks. 


SILICON   AND   BORON 


369 


The  alkali  silicates,  for  example,  sodium  silicate,  Na2SiO3,  can 
readily  be  obtained  by  fusing  silicon  dioxide  with  an  alkali  carbonate, 
and  are  easily  soluble  in  water.  If  hydrochloric  acid  is  added  to  a 
solution  of  sodium  silicate,  silicic  acid  js  thrown  down  as  a  gelatinous 
precipitate  ;  but  if,  on  the  other  hand,  a  dilute  solution  of  sodium 
silicate  is  cautiously  added  to  concentrated  hydrochloric  acid,  a  clear 
solution  containing  silicic  acid  is  obtained  : 

Na2SiOs+2HCl-»H2SiO3  +  2 


The  silicic  acid  is  not,  however,  in  true  solution,  but  in  what  is  termed 

colloidal  solution.     In  contrast  to  substances  present  in  true  solution, 

for  example,  sodium  chloride, 

colloids  do  not  pass  through 

certain  animal  and  vegetable 

membranes,  and  advantage 

may  be  taken  of  this  to  effect 

the  separation  of  the  silicic 

acid    and    sodium    chloride 

obtained  in  the  above  experi- 

ment.    The  arrangement  for 

this    purpose    is    shown    in 

Fig.  79.     Parchment  paper 

or  animal   bladder    is    tied 

tightly   over   the   end   of  a 

cylindrical    tube  ;   the   mix- 

ture containing  silicic  acid  is 

then  poured  into  the  vessel, 

which    is    suspended    in    a 

larger      vessel      containing 


FIG.  79. 


water,  as  shown.  The  sodium  chloride  and  excess  of  hydrochloric 
acid  readily  pass  through  the  membrane,  and  by  renewing  the 
water  occasionally  can  be  ultimately  almost  completely  removed  from 
the  solution.  The  clear  solution  of  silicic  acid  thus  obtained  is  fairly 
stable,  but  if  concentrated  beyond  a  certain  point  the  acid  separates 
in  a  gelatinous  form.  The  arrangement  just  described  is  termed  a 
dialyser,  and  the  process  is  called  dialysis. 

Silicic  acid  is  an  extremely  weak  acid,  as  is  shown  by  the  fact  that 
its  solutions  have  no  effect  upon  litmus  and  are  without  acid  taste, 
and  that  the  alkali  silicates  are  hydrolyzed  very  considerably  in 
solution. 

The  acid  H2SiO3  is  called  metasilicic  acid^  and  the  compound 
24 


370     A  TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

H4SiO4  orthosilicic  acid.  As  already  mentioned,  neither  has  been 
definitely  isolated,  and  the  silicic  acid  prepared  as  described  above 
doubtless  contains  a  mixture  of  acids. 

Silicates — As  already  indicated,  the  silicates  occurring  as  con- 
stituents of  rocks,  the  majority  of  which  are  well  crystallized,  may 
be  derived  from  hypothetical  silicic  acids  of  the  general  formula 
jrSiO2lyH2O.  From  metasilicic  acid,  H2SiO3,  are  derived  sodium  sili- 
cate, Na2SiO3,  and  ivollastonite,  Ca3Si3O9 ;  from  orthosilicic  acid, 
H4SiO4,  we  obtain  olivine,  Mg2SiO4.  Orthodase,  KAlSi3O8,  is  de- 
rived from  the  acid  H4Si3O8,  or  2H2O,3SiO2;  andalusite,  PA^iQ^ 
from  H6SiO6,  or  SiO2,3H2O  ;  and  meerschaum,  Mg2H4Si3O10,  from 
the  acid  H8Si3O10,  or  3SiO2,4H2O. 

Zeolites  are  hydrated  silicates,  for  example,  thomsonite, 
CaAl2Si2O8,3H2O.  They  have  the  remarkable  property  of  losing 
water  while  still  remaining  a  single  phase,  and  the  water  is  therefore 
not  present  as  water  of  crystallization,  but  is  simply  absorbed. 

The  silicates  of  which  the  original  rocks  of  the  earth's  crust  almost 
entirely  consist  are  continually  being  decomposed  by  water  and  the 
carbon  dioxide  of  the  atmosphere.  Perhaps  the  most  powerful  disinte- 
grating influence  is  due  to  alternate  freezing  and  thawing.  Moisture 
penetrates  into  small  cracks,  and  on  solidification  exerts  enormous  pres- 
sure. Carbon  dioxide,  as  it  forms  with  water  a  stronger  acid  than 
silicic  acid,  displaces  the  latter  from  combination,  silica  being  set  free. 
Sand,  which  is  found  in  large  deposits,  is  formed  in  this  way,  the 
carbonates  of  the  metals  being  washed  into  the  soil,  from  which 
they  are  largely  taken  up  by  plants.  Aluminium  silicate,  which  along 
with  alkali  silicates  occurs  in  many  rocks,  happens  to  be  particularly 
stable,  and  when  the  minerals  are  decomposed  under  the  combined 
influence  of  moisture  and  carbon  dioxide,  the  alkali  carbonates  are 
washed  away  and  the  aluminium  silicate,  being  deposited  in  a  finely- 
divided  form,  constitutes  clay. 

In  the  laboratory  silicates  are  decomposed  by  fusing  with  excess  of 
alkali  carbonate.  Sodium  silicate  and  carbonates  of  the  bases  are 
usually  formed  ;  the  former  can  be  dissolved  out  with  water,  and  the 
latter  are  usually  soluble  in  hydrochloric  acid. 

Colloidal  Solutions  '—We  have  already  seen  that  a  solution  of 
silicic  acid  is  unable  to  pass  through  an  animal  or  vegetable  mem- 
brane, whereas  dissolved  sodium  chloride  readily  passes  through. 
These  two  substances  may  be  taken  as  types  of  two  great  groups  of 
compounds,  crystalloids  and  colloids.  The  great  majority  of  the  com- 
pounds so  far  considered  (salts,  acids,  and  bases)  belong  to  the  class 
1  For  fuller  treatment  see  Physical  Chemistry,  chap,  xii. 


SILICON    AND    BORON 


37* 


of  crystalloids.  They  diffuse  relatively  fast  in  solution,  are  able  to  pass 
through  animal  and  vegetable  membranes,  and  can  usually  be  obtained 
without  difficulty  in  crystalline  form.  Colloids,  on  the  other  hand, 
diffuse  very  slowly  in  solution,  are  unable  to  pass  through  mem- 
branes, and  when  separated  from  solution  have  a  gelatinous  appear- 
ance, and  cannot  be  obtained  in  crystalline  form.  Besides  silicic 
acid,  gelatine,  gum,  caramel  or  burnt  sugar  and  arsenious  sulphide 
are  typical  colloids.  As  diffusive  power  is  proportional  to  osmotic 
pressure  (p.  195),  the  above  facts  may  also  be  stated  in  the  form 
that  the  osmotic  pressure  of  colloidal  solutions  is  very  small,  from 
which  it  follows  at  once  that  colloids  have  very  high  molecular 
weights. 

The  properties  of  colloidal  solutions  may  be  studied  very  con- 
veniently with  a  solution  of  arsenious  sulphide,  which  is  obtained  by 
dissolving  arsenious  oxide  in  distilled  water  by  boiling,  and  then 
passing  hydrogen  sulphide  into  the  cooled  solution.  A  yellow  fluor- 
escent solution  is  thus  obtained,  which,  when  examined  by  transmitted 
light,  appears  homogeneous  even  under  the  microscope.  When 
examined  with  the  ultra-microscope,  however,  it  is  seen  that  the  liquid 
is  full  of  minute  particles  in  extremely  rapid  motion.  We  have  here 
an  indication  of  the  fundamental  difference  between  solutions  of 
crystalloids  and  of  colloids  ;  in  the  former  the  particles  of  the  solute 
are  of  molecular  dimensions,  and  entirely  beyond  the  range  of  any 
microscope.  In  colloidal  solutions,  on  the  other  hand,  the  particles 
are  much  larger,  and  in  some  cases  can  be  seen  with  the  ordinary 
microscope. 

Many  colloids,  e.g.  arsenious  sulphide,  are  at  once  precipitated  from 
solution  on  addition  of  a  small  quantity  of  an  electrolyte,  while  non- 
electrolytes  are  without  action.  The  precipitated  colloid  is  termed  a 
gelj  when  in  solution  it  is  called  a  sol.  When  water  is  the  solvent 
concerned,  the  terms  hydrogel  and  hydrosol  are  used.  Further,  when 
two  electrodes  are  immersed  in  a  colloidal  solution  and  a  large  differ- 
ence of  potential  is  established  between  them,  the  colloidal  particles 
move  either  towards  the  positive  or  the  negative  pole.  This  pheno- 
menon is  best  accounted  for  on  the  view  that  colloidal  particles,  like 
ions,  are  associated  with  electrical  changes,  either  positive  or  nega- 
tive. Arsenious  sulphide  and  silicic  acid  are  negatively  charged ; 
ferric  hydroxide  is  positively  charged. 

Some  colloids  when  separated  in  the  gelatinous  form  (hydrogel) 
again  go  into  colloidal  solution  (hydrosol)  on  treatment  with  water, 
while  others,  e.g.  ferric  hydroxide,  are  unaffected  by  further  treatment 


372     A   TEXT-BOOK    OF   INORGANIC    CHEMISTRY 

with  water.     The  former  are  termed  reversible,  the  latter  irreversible 
colloids. 

It  is  not  correct  to  assume  that  a  "  colloidal  solution  "  is  necessarily 
a  solution  of  a  substance  which  appears  under  all  circumstances  as  a 
colloid.  (The  colloidal  form  is  also  a  condition  which  many  substances 
usually  appearing  as  crystalloids,  e.g.  arsenious  sulphide,  platinum, 
may  be  made  to  assume.  The  typical  colloids,  however,  such  as 
gum,  gelatine  and  silicic  acid,  never  appear  in  other  than  the  col- 
loidal condition. 

BORON 

Symbol,  B.     Atomic  weight  =  n.o.     Molecular  weight  unknown. 

Chemical  Relationships  —  -Boron  acts  exclusively  as  a 
trivalent  element,  and  belongs  to  the  same  group  as  aluminium 
(p.  479).  It  shows  comparatively  little  analogy  to  aluminium,  and, 
apart  from  the  types  of  the  compounds,  it  shows  much  more  analogy 
with  silicon.  One  oxide  of  boron,  B2O3,  and  a  number  of  oxyacids 
derived  from  it  are  known. 

Occurrence  —  Boron  does  not  occur  free  in  nature.  In  the  form 
of  boric  acid,  H3BO3,  it  is  contained  in  the  jets  of  steam  —  so-called 
soffioni  —  escaping  from  the  earth  in  volcanic  regions  in  Italy.  An 
important  source  of  boron  compounds  is  the  crude  borax  or  tincal, 
Na2B4O7,ioH2O,  found  in  the  United  States  (chiefly  in  Nevada  and 
California)  and  in  Tibet.  Other  minerals  containing  boron  are 
colemanite,  2CaO,3B2O3,5H2O,  and  boracite,  6MgO,8B2O3,MgCl2. 
Certain  silicates,  e.g.  tourmaline^  also  contain  boron. 

Modes  of  Formation  —  The  methods  of  obtaining  boron  are 
closely  analogous  to  those  used  for  obtaining  silicon. 

(1)  By  heating  boron  trioxide  (or  fused  boric  acid)  with  metallic 
sodium,  potassium  or  powdered  magnesium  in  a  covered  crucible  : 

B2O3  +6Na->2B  +  3Na2O. 

The  product  is  treated  with  dilute  hydrochloric  acid,  separated  from 
the  solution  by  filtration,  and  finally  washed  with  water. 

(2)  By  passing  the  vapours  of  boron  trifluoride  or  trichloride  over 
heated  metallic  sodium  : 

BC1 


(3)  By  decomposing  an  alkali  borofluoride  with  an  alkali  metal  or 
with  magnesium  : 

NaBF4+  3Na->B  +  4NaF. 


SILICON   AND    BORON  373 

(4)  The  boron  obtained  by  the  methods  already  described  is  the 
amorphous  variety.  Crystalline  boron  can  be  prepared  by  heating 
together  the  amorphous  modification  and  metallic  aluminium  at  1500° 
in  absence  of  air  for  i  to  2  hours,  cooling,  and  dissolving  out  the 
aluminium  with  hydrochloric  acid. 

Properties  —  (a)  Amorphous  Boron  —  This  modification  forms 
a  brown  to  brownish-black  powder  of  density  2.45  ;  it  does  not  fuse 
even  at  the  temperature  of  the  electric  arc.  It  combines  directly 
with  chlorine  at  410°,  with  sulphur  at  610°,  with  oxygen  at  700°  and 
with  nitrogen  at  1000°.  In  the  latter  case  boron  nitride,  BN,  is 
obtained,  and  therefore  when  boron  burns  in  air  a  mixture  of  oxide 
and  nitride  is  formed. 

Boron  is  not  affected  by  water  or  by  hydrochloric  acid,  but  is 
oxidized  on  heating  with  nitric  or  with  concentrated  sulphuric  acid  : 

2B+6HNO3->B2O3+6NO2+3H2O. 
On  fusing  with  alkalis  or  alkali  carbonates  borates  are  obtained  : 


(b)  "  Crystalline  Boron"  —  One  substance  described  under  this 
name  occurs  in  black  crystals,  and  appears  to  be  a  compound  of 
aluminium  and  boron,  A1B12.  It  is  nearly  as  hard  as  the  diamond. 
A  second  impure  modification,  occurring  in  yellow  transparent 
crystals,  has  also  been  described.  Up  to  the  present  crystalline 
boron  has  not  been  obtained  pure. 

Boron  Hydrides  —  When  magnesium  boride,  Mg3B2,  obtained  by  the  action  of 
magnesium  on  boron  trioxide,  is  treated  with  hydrochloric  acid,  the  mixture  of  gases 
condensed  in  a  U  tube  immersed  in  liquid  air,  and  the  mixture  then  fractionated, 
two  hydrides,  B4H10  (the  main  product)  and  B6H12,  are  obtained. 

The  hydride  B4H10  is  a  colourless  liquid,  b.pt.  i6-i7°/76o  mms.  spontaneously 
inflammable  in  air,  disagreeable  odour,  decomposes  in  a  few  hours  at  room 
temperature  forming  other  hydrides,  slowly  decomposed  by  water  and  by  alkali  : 
B4H10+  i2H20->4B(OH)3+  nH2. 

The  hydride  B6H12  is  also  a  colourless  liquid,  b.pt.  ioo°/76o  mms.  spontaneously 
inflammable,  more  sensitive  to  moisture  than  B4HW. 

The  hydride  B2H6,  formed  when  B4H10  is  kept  over  mercury  at  room  tempera- 
ture, is  a  colourless  gas  with  a  repulsive  odour,  not  spontaneously  inflammable  in 
air  when  pure,  decomposed  quantitatively  by  water;  B2H6+H2O-»2H3BO3+6H2. 

The  hydride  B10H14,  obtained  by  heating  B4H10  for  4-5  hours  at  100°,  is  obtained 
by  sublimation  in  colourless  needles,  m.p.99.5°  ;  fairly  stable  in  air,  not  attacked 
even  by  boiling  water. 

From  the  formula  B2H¥  it  appears  that  the  maximum  valency  of  boron  is  at 
least  four. 


374     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

Boron  Trifluoride,  BF3,  is  obtained  by  heating  an  intimate 
mixture  of  calciuhi  fluoride  and  boron  trioxide  : 

3CaF2  +  2B2O3->Ca3(BO3)2+2BF3, 

or  by  heating  together  boron  trioxide,  sulphuric  acid,  and  calcium 
fluoride  : 

(i.)  CaF2  +  H2SO4-»CaSO4 

(ii.)  B2 


Properties  —  Boron  trifluoride  is  a  colourless  gas  with  a  suffocat- 
ing odour.  It  fumes  in  the  air,  and  is  readily  taken  up  by  water, 
with  formation  of  hydrofluoboric'acid,  HBF4,  and  boric  acid  ;  the 
latter,  on  account  of  its  limited  solubility  in  water,  partially  crystal- 
lizes out  : 


Hydrofluoboric  add,  HBF4,  may  be  regarded  as  being  formed  by  the 
association  of  one  molecule  each  of  the  trifluoride  and  hydrofluoric 
acid,  BF3,HF,  and  in  fact,  is  partially  decomposed  into  its  com- 
ponents in  aqueous  solution  : 

HBF4^BF3+HF. 

Hydrofluoboric  acid  is  only  known  in  aqueous  solution.  Its  salts,  the 
ftuoborates,  are  stable.  The  whole  behaviour  of  boron  trifluoride 
strongly  recalls  that  of  silicon  tetrafluoride. 

Boron  Trichloride,  BC13—  Preparation—  (i)  By  direct  com- 
bination of  the  elements  above  400°. 

(2)  By  strongly  heating  a  mixture  of  boron  trioxide  and  charcoal 
and  passing  chlorine  over  it  : 


Properties  —  Boron  trichloride  is  a  colourless,  mobile  liquid,  which 
boils  at  1  8°.  It  fumes  in  the  air,  and  in  conformity  with  the 
general  behaviour  of  the  chlorides  of  non-metals  is  immediately 
decomposed  by  water,  hydrochloric  acid  and  boric  acid  being 
formed  : 

BC13  +  3HOH-»B(OH)3+3HC1. 

Boron  Nitride,  BN,  is  obtained  by  direct  combination  of  its 
components  at  a  white  heat,  or  by  strongly  heating  amorphous  boron 
in  ammonia  gas  or  in  nitric  oxide.  It  forms  a  white  amorphous 
powder,  which  is  stable  on  heating  in  air,  but  is  decomposed  when 


SILICON    AND    BORON  375 

heated  with  steam  at  200°  under  pressure,  ammonia  and  boric  acid 
being  formed  : 

BN  +  3HOH->NH3+H3BO3. 

It  is  also  decomposed  by  fusing  with  potassium  hydroxide  : 


Boron  Trisulphide,  B2S3,  is  obtained  by  direct  combination  of 
its  elements  or  by  heating  a  mixture  of  boron  trioxide  and  carbon  in 
a  current  of  carbon  disulphide.  It  occurs  in  small,  colourless  needles, 
and  is  immediately  decomposed  by  water,  with  formation  of  boric 
acid  and  hydrogen  sulphide  : 

B2S3-f6H2O-»2H3 


OXIDES  AND  OXYACIDS  OF  BORON 

As  already  indicated,  only  one  oxide  of  boron,  B2O3,  is  known. 
Corresponding  with  this  oxide  are  the  well-defined  oxyacids  : 

Orthoboric  acid         ....     H3BO3     HO-B:(OH)2 
Pyroboric  acid  .....     H2B4O7 
Metaboric  acid  .....     HBO2     HO-B  =  O 

Boron  Trioxide,  B2O3,  is  formed  when  boron  is  burned  in  air 
or  oxygen,  and  is  usually  prepared  by  heating  boric  acid  to  redness 
in  a  crucible  : 


Properties  —  Boric  acid  is  a  transparent,  glassy,  hygroscopic 
solid,-  which  melts  about  580°,  and  can  be  raised  to  a  very  high 
temperature  without  appreciable  volatilization.  On  this  account  it 
displaces  volatile  acids  from  their  salts  at  a  •  high  temperature  ; 
carbonates  and  nitrates  are  completely,  sulphates  only  partially, 
decomposed  : 


It  combines  with  many  metallic  oxides  when  fused,  giving  glasses  ol 
characteristic  colours  ;  this  fact  is  taken  advantage  of  in  testing  for 
certain  metals. 

Orthoboric    Acid,    H3BO3  —  Preparation  —  (i)  On  the  com- 
mercial scale,  boric  acid  is  usually  obtained  from  the  "  soffioni  "  in 


376     A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 


,  in  which  it  occurs  to  a  very  small  extent  (less  than  o.i  per 
cent.).  The  jets  are  caused  to  pass  through  water  contained  in  brick 
reservoirs  built  round  them,  and  after  the  water  has  been  passed 
through  several  reservoirs,  and  has  thus  taken  up  a  considerable 
proportion  of  the  acid  and  become  considerably  concentrated  by 
evaporation,  it  is  further  evaporated  by  the  heat  of  the  escaping 
vapours  and  the  solution  then  set  aside  to  crystallize. 

The  source  of  the  boric  acid  in  the  jets  is  not  thoroughly  under- 
stood. It  probably  results  from  the  action  of  steam  on  boron  nitride 
or  on  boron  sulphide  (q-v.}  below  the  surface  of  the  earth. 

(2)  Boric  acid  is  also  prepared  by  adding  sulphuric  or  hydrochloric 
acid  to  a  concentrated  solution  of  borax  (sodium  pyroborate, 
Na,B4O7)  : 


Properties  —  Boric  acid  occurs  in  colourless,  lustrous,  transparent 
six-sided  leaflets,  which  are  soapy  to  the  touch.  It  is  moderately 
soluble  in  water.  At  13°  100  c.c.  of  the  saturated  solution  contain 
3.84  grams,  at  20°  4.91  grams,  and  at  25°  5.58  grains  of  the  acid. 
When  the  aqueous  solution  is  boiled  a  little  boric  acid  escapes  with 
the  steam.  The  acid  is  readily  soluble  in  alcohol,  and  when  the 
solution  is  ignited  it  burns  with  a  green-edged  flame.  This  property 
is  used  as  a  test  for  the  acid  and  its  salts. 

Boric  acid  is  an  extremely  weak  acid.  The  aqueous  solution 
colours  litmus  a  sort  of  port  wine  colour,  instead  of  the  bright  red 
produced  by  strong  acids.  Even  in  the  first  stage  of  the  dissociation 


the  ionisation  is  much  less  than  for  carbonic  acid,  and  it  is  therefore 
impossible,  owing  to  hydrolysis,  to  obtain  corresponding  salts  — 
orthoborates  —  from  aqueous  solution. 

When  boric  acid  is  heated  for  some  time  at  Jtoq^,  it  loses  water, 
with  formation  of  metaboric  acid:  H3BO3->HBO2  +  H2O,  and  at 
140°  still  more  water  is  driven  off,  pyroboric  acid  being  formed  : 
4HBO2->H2B4O7  +  H2O. 

Metaboric  Acid,  HBO2,  is  obtained  as  already  mentioned,  by 
heating  orthoboric  acid  to  ^ipo^  It  is  a  monobasic  acid,  and  a 
number  of  corresponding  salts,  e.g.  NaBO2,  AgBO2,  Ca(BO2)2  are 
known. 

Pyroboric  Acid,  H2B4Or,  is  obtained  by  heating  orthoboric 
acid  to  Jjo°.  When  pyroboric  acid  or  metaboric  acid  is  dissolved 
in  water,  orthoboric  acid  is  immediately  formed. 


SILICON  AND  BORON  377 

Boratas  —  The  best  known  borate  is  ordinary  borax,  sodium 
pyroborate,  Na2B4O7,ioH2O.  It  occurs  naturally  in  tincal  (p.  372), 
and  is  also  obtained  commercially  by  the  action  of  sodium  carbonate 
on  colemanite  : 

Ca2B6On  +  2Na2CO3->Na2B4O7  +  2  NaBO2  +  2CaCOs. 

The  solution  is  filtered  to  remove  the  calcium  carbonate,  and  on 
evaporation  the  borax  crystallizes  out.  The  more  soluble  metaborate 
remaining  in  the  mother  liquor  can  also  be  converted  into  borax  by 
passing  carbon  dioxide  through  the  solution  : 

+  CO2->Na2CO3+Na2B4Or. 


On  the  laboratory  scale,  borax  can  be  prepared  by  adding  sodium 
carbonate  to  a  boiling  solution  of  boric  acid,  evaporating  the  solu- 
tion; and  setting  aside  to  crystallize  : 


Borax  occurs  in  large  colourless  prisms.  On  heating  it  loses  water, 
swells  up,  and  finally  fuses  to  a  clear  bead,  to  which  characteristic 
colours  are  imparted  by  certain  metallic  oxides,  such  as  those  of 
cobalt  and  of  manganese.  If  the  formula  is  written  thus,  Na2O,2B2O3, 
we  can  understand  that  certain  basic  oxides  can  combine  with  part 
of  the  boron  trioxide,  forming  definite  chemical  compounds. 

Borax  is  hydrolyzed  to  a  considerable  extent  by  water,  and  the 
aqueous  solution  has  therefore  an  alkaline  reaction. 

As  already  indicated,  metaborates  have  also  been  obtained  from 
aqueous  solution  ;  thus  silver  metaborate,  AgBO2,  and  calcium  meta- 
borate, Ca(BO2)2,2H2O,  have  been  prepared  in  this  way. 

The  weakness  of  the  boric  acids  is  strikingly  shown  by  the  fact 
that  although  borax  contains  a  large  excess  of  the  acidic  oxide, 
Na2O  :  2B2O3,  its  aqueous  solutioji_Js_alkaline.  Moreover,  there  is 
evidence  that,  as  is  to  be  anticipated,  pyroboric  acid  is  a  consider- 
ably stronger  acid  than  orthoboric  acid. 


CHAPTER  XXV 

CLASSIFICATION  OF  THE  ELEMENTS— THE  PERIODIC 
SYSTEM— GENERAL  PROPERTIES  OF  THE  METALS 
AND  THEIR  COMPOUNDS 

HAVING  regard  to  the  enormous  number  of  facts  already 
established  with  reference  to  the  behaviour  of  the  elements  and 
their  compounds,  it  is  evident  that  some  system  of  classifying  the 
elements  which  will  show  their  mutual  relationships  and  facilitate 
the  comprehension  and  utilization  of  the  available  data  is  in  the 
highest  degree  desirable. 

Several  methods  of  classification  might  be  suggested.  For  ex- 
ample, the  elements  might  be  arranged  according  to  their  valencies, 
all  univalent  elements,  for  instance,  being  brought  into  one  group. 
There  are  two  serious  drawbacks  to  this  suggestion.  In  the  first 
place,  elements  like  potassium  and  iodine,  which  clearly  have  not 
the  slightest  analogy,  would  be  brought  together.  In  the  second 
place,  most  elements  have  more  than  one  valency,  and  might  there- 
fore with  equal  justification  be  put  in  two  or  more  groups. 

The  division  of  the  elements  into  metals  and  non-metals,  already 
made  use  of,  is  much  more  promising,  although  by  no  means 
satisfactory.  The  main  differences  between  the  two  groups  will  be 
referred  to  in  detail  later,  and  be  indicated  here  only  very  briefly. 
The  chief  characteristics  of  metals  are  (i)  metallic  lustre  and  ability 
to  conduct  heat  and  electricity  ;  (2)  they  combine  with  oxygen  to 
form  basic  oxides  ;  (3)  when  forming  constituents  of  salts  in  solution, 
they  appear  alone  as  positive  ions  only  ;  (4)  their  chlorides  (and 
other  binary  halogen  compounds)  are  fairly  stable  towards  water- 
In  contrast  to  the  metals,  the  typical  non-metals  (i)  have  no  metallic 
lustre  and  do  not  conduct  heat  and  electricity  ;  (2)  form  acidic 
oxides ;  (3)  appear  free  in  the  ionised  condition  as  negative  ions 
only  ;  and  finally,  (4)  their  compounds  with  the  halogens  are  readily 
hydrolyzed  by  water.  It  has  already  been  pointed  out,  and  further 
illustrations  will  be  given  later,  that  although  these  broad  differences 
are  quite  sufficient  to  distinguish  between  typical  metals  and  non- 
378 


CLASSIFICATION    OF   THE   ELEMENTS          379 

metals,  a  number  of  elements  appear  to  stand  on  the  border-line 
between  these  two  great  groups. 

By  far  the  most  successful  system  yet  proposed  for  classifying  the 
elements  is  based  upon  arranging  them  according  to  their  atomic 
weights.  In  1829,  Dobereiner  pointed  out  that  there  are  a  number 
of  groups  of  three  elements,  so-called  triads,  the  members  of  which 
have  close  chemical  analogy,  while  the  atomic  weights  differ  by  about 
16  or  a  multiple  of  that  number.  Examples — chlorine,  bromine  and 
iodine  ;  calcium,  strontium  and  barium.  In  1864,  Newlands  showed 
that  when  the  elements  then  known  were  arranged  in  the  order  of 
increasing  atomic  weights,  although  the  successive  elements  showed 
no  particular  analogy,  the  eighth  element  was  analogous  to  the  first, 
the  ninth  to  the  second,  and  so  on  ;  in  other  words,  when  the  elements 
are  arranged  as  above  stated,  there  is  a  periodic  recurrence  of  elements 
with  similar  chemical  properties.  As  each  such  group  or  period  con- 
tained seven  elements,  the  discovery  of  Newlands  was  termed  the 
law  of  octaves. 

In  1869,  Mendeleeff,  on  the  same  basis,  but  without  any  knowledge 
of  Newlands'  work,  developed  the  periodic  system  of  the  elements 
substantially  in  the  form  now  accepted.  Starting  with  helium =4, 
which  was  unknown  when  Mendeleeff  first  brought  forward  his  classi- 
fication, we  have  the  arrangement  shown  in  the  first  line  of  the 
accompanying  small  table— 

He=  4    Li  =  7    Gl  =  9     B  =11     C=i2     N  =  i4     O=i6    F  =19 
Ne=2o    Na=23    Mg=24    Al=27    81  =  28     P  =31     8=32    0=35.5 

in  which  the  properties  vary  regularly  from  helium  to  fluorine.  The 
next  element,  neon  =  20,  is  an  inactive  gas  and  is  placed  below  helium, 
sodium  then  falls  into  its  proper  place  below  lithium,  phosphorus 
below  nitrogen  (p.  238),  sulphur  below  oxygen  (p.  322).  It  should  be 
clearly  realized  that  no  elements  are  omitted  in  this  arrangement;  the 
order  is  strictly  that  of  increasing  atomic  weight,  and  the  accuracy 
with  which  the  elements  fall  into  the  places  to  which  they  would  be 
assigned  on  purely  chemical  grounds  is  most  striking.  These  groups 
of  elements  are  termed  periods,  and  the  first  two  periods  contain 
8  elements  each. 

The  complete  table,  arranged  on  the  above  principle,  is  given  on 
p.  380.  A  new  period  is  started  with  argon,  a  third  inactive  gas,  but 
in  this  case  it  is  necessary  to  pass  over  18  elements  before  another 
inactive  gas  (krypton),  bearing  a  strong  resemblance  to  argon,  is 
reached.  Such  a  period  of  18  elements  is  termed  a  long  period  in 


tN 

s 

II 

0 

II 

1 

a 

\  5 

1 

1 

*0\ 

J!        ! 

0? 

| 

w 

IT 

3 

|| 

cS 

'ft 

tN 

1 

| 

H 

1 

§, 

d 

' 

vn 

5 

$                           \ 

1 

HH  O 

^ 

10 

H 

~ 

|g 

II 

CO 

ITffl 

jl 

1 

w 

p 

c 

1               1 

1 

ON 

VO               ^O     | 

1 

so" 

M 

CO 

%«                ^ 

on 

CO 

> 

wS 

0 

II 
03 

1?^ 

Si!     i 

II 

II 

U 

S  H               £n 

^ 

D 

in 

S           »o  I 

10 

CO    US 

^, 

^ 

M                      0 

co  oo 

CO 

II 

$ 

5  ><?> 

II 

II 

s>^ 

ON  11                H 

H    "I 

1 

M  W 

^ 

M-< 

II 

jlco          H 

> 

S           c^1 

hffl 

so- 

n 

H 

od 
O 

00^ 

VO     M                    0      ' 

8 

II 

»0 

~ 

wS 

II 

U 

It 

o3 

II   v 
P° 

lfc/5         Jl 

N                U 

£ 
1 

II 
S 

Cfo 

H 

fr 

a 

a      cJ 

CON 

~ 

3  SI 

II 

CQ 

II 

n"o 

P      .3 

r 

1 

^ 

.'* 

J        I 

8 

vO 

09 

ON 

N 

0.^ 

o 

H 

d 

M 

" 

a« 

O 

II 

b/i 

5s 

"-S      "I 

03°            ^ 

if 

ej 

vO 

00 

tx 

'"'O 

tx 

o 

CO 

!£  s      co  i 

« 

sS 

II 

II 

?§ 

II  ho          |f 

II 

1 

* 

^U 

^  ^           (/) 
Di                U 

1  < 

J* 

0 

o 

O                        (N 
00                        H 

1 

T 

0 

w 

« 

II 

jj 

II 

3 

II          II 

1 

n 

ci, 

o 

c 

V 

ist  Short 
Period 

ll 
p 

y 

be                bo 
§|            §| 

n                co 

bo 

,-  ^ 
^- 

OJ3 

ll 

^ 

CLASSIFICATION    OF   THE    ELEMENTS         381 

contrast  to  the  two  short  periods  of  8  elements  each.  The  whole 
table  is  made  up  of  two  short  and  five  long  periods,  but  four  of  the 
long  periods  are  incomplete,  and  the  last  one  contains  only  3  elements. 
A  word  must  be  said  as  to  the  significance  of  the  blanks  in  the  table. 
After  molybdenum  =96  the  next  element  is  ruthenium  =  101.7,  which, 
strictly  speaking,  should  come  below  manganese  =55.  The  essential 
feature  of  this  arrangement,  however,  is  the  chemical  similarity  of 
elements  in  the  same  vertical  rows.  Now,  ruthenium  shows  no 
analogy  whatever  with  manganese,  but  shows  certain  analogies  with 
iron  =56.  It  is  therefore  placed  under  the  latter  element,  and  a  blank 
is  left  below  manganese,  to  indicate  the  position  of  a  hitherto  undis- 
covered element.  At  first  sight,  this  policy  seems  a  somewhat  arbitrary 
one,  but  it  has  been  entirely  justified  by  results  (see  below). 

The  arrangement  adopted  in  representing  the  long  periods  will  be 
evident  from  the  table.  The  twelfth  element  shows  some,  but  not  a 
very  close  analogy  to  the  second,  the  thirteenth  is  distantly  related 
to  the  third,  and  so  on.  The  former  element  of  the  pair  is  therefore 
placed  below,  but  to  the  right  of  the  latter,  to  indicate  the  absence  of 
a  close  analogy.  This  system  leaves  3  elements  in  the  middle  of  the 
first,  second  and  fourth  long  periods,"  the  position  of  which  presents 
some  difficulty.  Mendeleeff  put  these  "transition"  elements  in  a 
group  by  themselves,  the  so-called  eighth  group  (Group  vin.  in  the 
table). 

A  study  of  the  elements,  arranged  as  above,  shows  many  striking 
regularities.  Thus  the  valency  with  regard  to  hydrogen  increases 
regularly  up  to  the  middle  of  a  short  period,  and  then  falls  regularly 
(ultimately  to  unity)  in  the  latter  half  of  the  period,  whilst  the  maximum 
valency  for  oxygen  increases  regularly  by  units  from  the  beginning  up 
to  the  end  of  the  period.  The  valencies  in  the  long  periods  are  not 
quite  so  regular,  being  complicated  by  the  fact  that  most  elements  have 
several  valencies,  but,  generally  speaking,  each  of  the  halves  of  a  long 
period  behaves  like  a  short  period.  It  follows  at  once  from  this  arrange- 
ment that  the  elements  in  the  same  vertical  group  should  have  the 
same  valency,  and  a  glance  at  the  table  shows  that  such  is  the  case. 
The  group  under  o  comprises  the  inactive  gases,  which  do  not  enter 
into  chemical  combination,  and  may  therefore  fittingly  be  regarded 
as  having  zero  valency.  The  next  group,  under  I.,  composed  of 
univalent  elements,  comprises  the  alkali  metals,  which  are  typical 
univalent  elements,  and  copper,  silver  and  gold,  all  of  which  can  act 
as  univalent  elements.  Similar  considerations  apply  to  the  succeed- 
ing groups.  The  position  of  the  halogens  in  Group  VH.  is  interesting, 


382     A   TEXT  BOOK    OF   INORGANIC   CHEMISTRY 

and  appears  to  be  justified  to  some  extent  by  the  existence  of  com- 
pounds of  the  type  HC1O4,KIO4,  etc.  (p.  182).  The  members  of  the 
eighth  group  would  be  expected  to  show  a  valency  of  8.  This  is  to 
some  extent  borne  out  by  the  existence  of  osmium  tetroxide,  OsO4, 
and  ruthenium  tetroxide,  RuO4  (p.  572),  but  the  other  elements  do  not 
appear  to  exert  so  high  a  valency. 

Many  of  the  physical  properties  of  the  elements,  such  as  the 
melting-point,  the  atomic  volume,  the  density,  conductivity  of  the 
metals  for  heat  and  electricity,  heat  of  formation  of  oxides  and 
chlorides,  also  vary  regularly  within  such  period.  For  example,  the 
melting-points  of  the  members  of  the  first  period  gradually  rise  from 
helium  to  carbon  and  fall  again  to  fluorine,  and,  similarly,  the  elements 
of  highest  melting-point  (iron,  cobalt,  nickel,  etc.)  occur  in  the  middle 
of  the  long  periods. 

Not  only  the  physical  properties,  but  also  the  chemical  properties 
of  the  elements  vary  regularly  within  each  period.  Thus  the  elements 
on  the  extreme  left  hand  are  inactive  gases,  those  in  the  second  group 
(Series  A)  decompose  water  and  are  strongly  electro-positive,  at  the 
middle  of  the  period  the  electrical  character  is  much  less  pronounced 
(carbon,  silicon),  and  towards  the  right  hand  strongly  electro-negative 
elements  are  found. 

Besides  this  variation  of  properties  in  the  horizontal  series,  there  is 
a  similar,  but /much  less  marked  variation  of  both  physical  and 
chemical  properties  in  the  vertical  groups.  In  the  case  of  the  alkali 
metals,  for  instance,  the  melting-points  fall  from  lithium  to  caesium, 
and  the  electro-positive  character  increases. 

The  statements  in  the  last  three  paragraphs  are  summarized  in 
the  Periodic  Law,  due  to  Mendeleeff,  which  may  be  stated  as 
follows  : 

The  properties  (both  physical  and  chemical)  oj  the  elements  and 
their  compounds  are  periodic  functions  of  the  atomic  weights. 

The  Periodicity  of  Physical  Properties — It  appears 
desirable  to  illustrate  the  periodicity  of  the  physical  properties  of  the 
elements  a  little  more  in  detail. 

(a)  The  Density  of  the  Elements — The  density  increases  regularly 
from  the  beginning  to  the  middle  of  a  period,  short  or  long,  and  then 
diminishes  regularly  towards  the  end  of  the  period.  Uncertainty 
often  arises  owing  to  the  existence  of  a  number  of  forms  of  some 
elements  with  different  densities,  and  also  as  to  the  proper  conditions 
for  comparison.  This  rule  is  illustrated  in  the  accompanying  small 
table  from  the  data  for  the  second  short  period  : 


CLASSIFICATION    OF   THE    ELEMENTS         383 

Element      Na         Mg          Al         Si      P  (red)       S  Cl 

(rhombic)  (liquid) 
Density       0.97        1.75       2.67       2.49       2.14       2.06         1.33 

The  same  periodic  variation  occurs  in  the  density  of  the  oxides  of 
these  elements : 

Oxide         Na2O     MgO      A12O3      SiO2      P2O6         SO3      C12O7 
Density        2.8          3.7          4.0          2.6          2.7          1.9 

(b)  The  Atomic  Volumes  of  the  Elements — In  the  same  way  the 
atomic  volumes  of  the  elements — that  is,  the  relative  volumes  occupied 
by  the  atomic  weight  in  grams  of  the  elements  in  the  solid  state — 
vary  in  a  periodic  manner,  but  in  this  instance  the  magnitude  dimin- 
ishes towards  the  middle  of  a  period  and  then   increases  regularly 
towards  the  end.     The  elements  of  smallest  atomic  volume  therefore 
occur  towards  the  middle  of  a  period,  whether  short  or  long  (examples  ; 
boron,  aluminium,  iron,  cobalt,  nickel)  and  those  of  largest  atomic 
volume   at   the    ends    of  the    period.      These   facts   are   illustrated 
graphically   in   the  accompanying  diagram   (Fig.  80),  in  which  the 
atomic  volumes  as  ordinates  are  plotted  against  the  corresponding 
atomic  weights.     The  periodic  variation  of  this  property,  resembling 
the  successive  swings  of  a  pendulum  of  ever-increasing  amplitude,  is 
very  striking. 

(c)  The  Melting-points  of  the  Elements — The  melting-points  of  the 
elements  rise  towards  the  middle  of  a  period,  whether  short  or  long, 
and  fall  again  towards  the  end.     This  fact  is  illustrated  by  the  data 
for  the  elements  of  the  second  short  period  : 

Element  Ne    Na       Mg        Al  Si          PS        Cl 

Melting-point       -     97°     <8oo°    657°    >i7oo°    45°    115°    -102° 

The  results  in  this  instance  are  not  quite  so  regular,  but  the  general 
tendency  of-the  numbers  is  quite  definite. 

Uses  of  the  Periodic  System — The  periodic  system  is  of 
use  mainly  in  three  ways  : 

(1)  As  a  system  of  classification  which  indicates  in  a  fairly  satis- 
factory way  the  physical  and  chemical  relationships  of  the  elements. 

(2)  For  predicting  the  existence  and  properties  of  elements  hitherto 
undiscovered. 

(3)  For  enabling  us  to  find  the  correct  values  of  the  atomic  weights 
of  elements  which  do  not  form  volatile  compounds. 


384     A   TEXT-BOOK    OF   INORGANIC    CHEMISTRY 
o        o         o       /  o         o         o         o 

— <« oo  f*»  <O  iO      .        •<*•  rO 


CLASSIFICATION   OF   THE    ELEMENTS         385 

The  first  point  has  already  been  sufficiently  illustrated  in  the  fore- 
going paragraphs.  In  the  remainder  of  the  book,  the  periodic 
system  will  be  taken  as  the  basis  of  discussion  of  the  metals  and  their 
compounds,  and  its  great  value  will  then  become  apparent. 

As  regards  the  second  point,  when  the  periodic  system  was  first 
brought  forward,  there  were  more  blanks  in  the  table  than  there  are 
at  the  present  day,  and  Mendeleeff  not  only  suggested  that  the 
positions  of  these  blanks  corresponded  with  hitherto  undiscovered 
elements,  but  even  foretold  the  properties  of  the  missing  elements  from 
those  of  the  known  elements  near  them  in  the  periodic  table.  At 
that  time  (1871)  there  was  a  blank  in  the  table  representing  an 
unknown  element  with  an  atomic  weight  somewhat  exceeding  that  of 
zmc,  and  Mendeleeff,  from  the  properties  of  the  surrounding  known 
elements,  zinc  and  arsenic  in  the  horizontal  and  aluminium  and 
indium  in  the  vertical  rows,  foretold  that  the  unknown  element,  which 
he  termed  eka-aluminium,  would  have  an  atomic  weight  of  about  69, 
that  it  would  be  trivalent  and  form  alums  (p.  484)  like  aluminium, 
that  its  density  would  be  about  5.9,  that  it  would  be  more  volatile  than 
aluminium,  and  therefore  its  discovery  might  be  expected  by  means 
of  the  spectroscope.  Four  years  afterwards,  Lecoq  de  Boisbaudran 
was  led  by  spectroscopic  observations,  exactly  as  Mendeleeff  had 
foretold,  to  the  discovery  of  an  element,  which  he  called  gallium, 
having  all  the  properties  predicted  for  eka-aluminium.  Later,  two 
other  blanks  were  filled  up  by  the  discovery  of  scandium  (Nilson  and 
Cleve,  1879)  and  of  germanium  (Winkler,  1886),  which  also  showed 
in  all  respects  the  properties  predicted  by  Mendeleeff. 

The  use  of  the  periodic  system  for  fixing  atomic  weights  will  be 
readily  understood  from  the  foregoing.  When  the  equivalent  of  the 
element  has  been  determined,  it  is  usually  possible  to  decide  which 
multiple  of  it  is  to  be  taken,  as  there  will  in  general  be  only  one 
position  in  the  table  into  which  the  element  can  be  satisfactorily 
fitted.  The  most  familiar  example  is  the  controversy  with  regard  to 
the  atomic  weight  of  beryllium  (glucinum).  From  analysis  of  the 
chloride,  the  equivalent  of  this  element  was  found  to  be  4.5.  If  the 
metal  is  bivalent,  the  atomic  weight  must  be  9.0  and  the  formula  of  the 
chloride  is  BeCl2,  if,  on  the  other  hand,  it  is  trivalent,  the  atomic 
weight  is  13.5  and  the  formula  of  the  chloride  BeCl3.  Mendeleeff 
pointed  out  that  a  trivalent  element  of  atomic  weight  13.5  cannot  be 
fitted  into  the  periodic  table,  but  a  bivalent  element  of  atomic  weight 
9.0  fills  the  space  between  lithium  =  7.0  and  boron  =  u.o.  The  state- 
ment gave  rise  to  a  controversy  which  lasted  about  ten  years,  and 
25 


386     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

which  was  finally  settled  in  favour  of  Mendele'eff  by  Nilson  and 
Petterson,  two  of  the  chief  advocates  of  the  trivalency  of  beryllium. 
These  chemists  determined  the  vapour  density  of  beryllium  chloride 
to  be  41,  hence  the  molecular  weight  of  the  chloride  is  82,  and  its 
formula  BeCl2. 

Deficiencies  of  the  Periodic  System — Although  the 
periodic  system  has  proved  of  the  highest  value  for  the  development 
of  chemistry,  it  must  be  admitted  that  it  is  by  no  means  satisfactory 
in  all  respects.  We  shall  learn  later  that  there  is  not  that  close 
analogy  between  the  members  of  the  copper  group  (copper,  silver  and 
gold)  and  the  alkali  metals  to  be  anticipated  from  their  position 
in  the  first  group,  and  it  has  been  suggested  that  the  first  three 
elements  belong  more  properly  to  the  eighth  group,  coming  after 
nickel,  palladium  and  platinum  respectively. 

There  are  several  instances  in  which  the  order  of  the  atomic  weights 
is  different  from  that  required  by  the  chemical  behaviour  of  the 
elements.  Thus  argon  must  undoubtedly  come  below  neon,  and  be 
followed  by  potassium,  which  fits  into  its  proper  place  below  sodium, 
yet  the  atomic  weight  of  argon  is  greater  than  that  of  potassium. 
The  question  which  has  raised  most  discussion  in  this  connexion, 
however,  is  the  relative  position  of  tellurium  and  iodine.  Although 
from  its  chemical  relationships  (p.  321)  the  latter  element  must  follow 
iodine,  yet  experiment  shows  that  the  atomic  weight  of  tellurium  is 
greater  than  that  of  iodine.  It  is  at  first  natural  to  suppose  that  there 
must  have  been  some  error  in  determining  the  atomic  weights,  but 
recent  investigations  appear  definitely  to  have  disproved  this  sugges- 
tion. Finally,  the  chemical  behaviour  of  cobalt  indicates  that  it 
should  precede  nickel  in  the  periodic  table,  whereas  its  atomic  weight 
is  somewhat  greater  than  that  of  the  latter  element. 

There  is  no  definite  place  for  hydrogen  in  the  table.  Ramsay 
places  it  at  the  summit  of  the  halogen  group,  but  it  must  be  admitted 
that  the  analogies  are  slight.  Other  chemists,  with  perhaps  even 
less  justification,  place  hydrogen  at  the  summit  of  the  alkali  metals ; 
recent  investigation  shows  that  it  has  no  metallic  properties. 

The  position  of  the  metals  of  the  rare  earths  in  the  periodic  system 
is  very  uncertain  (p.  490). 

These  considerations  show  that  the  periodic  system,  although  of 
the  highest  value,  is  only  a  first  approximation  to  a  satisfactory 
system.  The  difficulties  of  classification  arise  from  the  many-sided 
character  of  the  different  elements,  as  illustrated  in  their  physical  and 
more  particularly  in  their  chemical  properties. 


GENERAL   PROPERTIES   OF   THE   METALS     387 


GENERAL  PROPERTIES  OF  THE  METALS 

It  will  be  of  advantage,  before  taking  up  the  study  of  the  metals 
and  their  compounds  in  detail,  to  give  a  brief  account  of  their  general 
characters,  more  particularly  as  to  the  methods  of  preparation  and 
general  properties  of  the  salts. 

It  has  already  been  stated  that  the  metals  are  broadly  distinguished 
from  the  non-metals  (i)  by  their  metallic  lustre  and  conducting 
power  for  heat  and  electricity  ;  (2)  by  uniting  with  oxygen  to  form 
basic  oxides,  which  latter  unite  with  acids  to  form  salts  ;  (3)  the 
metal  forming  a  constituent  of  a  salt  appears  alone  in  the  form  of 
positive  ions  only ;  (4)  the  chlorides  are  fairly  stable  towards  water. 
A  few  general  remarks  on  these  characters  will  be  of  interest. 

The  conducting  power  of  metals  for  electricity  is  expressed  in  terms 
of  the  conductivity,  in  reciprocal  ohms,  of  a  cm.  cube  of  the  substance  ; 
this  magnitude  is  termed  the  specific  conductivity,  k.  The  metal  of 
highest  conductivity  is  silver,  and  copper  comes  very  near  to  it  in 
this  respect.  Some  non-metals,  in  one  or  other  of  their  allotropic 
modifications  (e.g.  carbon)  show  conducting  power,  but  are  much 
less  efficient  conductors  than  the  metals.  Solutions  which  show 
electrolytic  conduction  are  all  far  inferior  to  the  metals  in  this  regard. 
The  following  small  table,  which  illustrates  these  statements,  is  in- 
structive ;  the  numbers  are  valid  for  18°. 

Substance  Silver       Copper    Mercury     Gas          30  per  cent. 

Carbon    Sulphuric  Acid 
Sp.  conductivity  624,000     587,000      10,240       200  1.35. 

The  points  (2)  and  (3)  are  closely  allied.  The  majority  of  basic 
oxides  combine  more  or  less  readily  with  water  to  form  hydroxides, 
which  in  solution  are  electrolytically  dissociated  into  positive  metallic 
ions  and  negative  OH'  ions.  Hydroxides  which  behave  in  this  way 
are  termed  bases,  and  are  characterized  by  the  properties  already 
fully  described.  As  regards  those  metallic  oxides  which  are  practi- 
cally insoluble  in  water,  e.g.  magnesium  oxide,  their  basic  character 
is  recognized  by  the  fact  that  with  acids  they  form  salts,  whose 
positive  ions  in  solution  consist  of  charged  atoms  of  the  metal. 

The  "  strength "  of  a  base  is  measured,  as  already  stated,  by  the 
extent  to  which  it  is  dissociated  in  solution,  but  the  application  of 
this  criterion  to  metallic  hydroxides  is  often  complicated  by  the 
slight  solubility  of  the  latter.  Another  method  of  measuring  the 


388     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

strength  of  a  base  is  to  determine  the  extent  to  which  a  salt  is 
hydrolyzed  in  aqueous  solution.  It  is  evident  that  if  we  compare 
in  this  respect  salts  formed  by  different  metals  with  the  same  strong 
acid,  for  example  hydrochloric  acid,  the  salts  of  strong  bases  will 
undergo  little  or  no  hydrolysis,  but  as  the  basic  character  diminishes 
the  hydrolysis  will  become  increasingly  pronounced  (cf.  p.  267). 
These  considerations  enable  us  to  understand  criterion  (4)  for  the 
metals ;  when  the  basic  character,  in  other  words  the  tendency  to 
form  positive  ions,  is  practically  or  entirely  absent  the  chloride  is 
readily  decomposed  by  water  (e.g.  phosphorus  trichloride). 

The  hydroxides  of  non-metals,  e.g.  P(OH)3  or  H3PO4,  have  a 
tendency  to  form  salts  with  bases  ;  in  other  words,  the  hydroxides 
are  acidic.  We  shall  learn  later  that  the  hydroxides  of  certain  metals 
have  both  basic  and  acidic  properties. 

Another  way  of  regarding  the  matter  is  to  consider  the  relative 
affinity  for  electricity  shown  by  different  elements  in  the  atomic  con- 
dition. Metals  which  have  a  great  tendency  to  form  positive  ions 
may  be  regarded  as  having  a  great  affinity  for  positive  electricity  ;  in 
other  words,  such  metals  are  strongly  electro-positive  (cf.  p.  434).  The 
non-metals  have  no  affinity  for  positive  electricity  ;  on  the  contrary, 
some  of  them,  such  as  the  halogens,  have  a  great  affinity  for  negative 
electricity,  and  are  therefore  said  to  be  strongly  electro-negative 
elements.  They  form  negative  ions  in  solution.  Finally,  certain 
elements,  such  as  carbon  and  silicon,  appear  to  have  little  or  no 
tendency  to  associate  with  electricity ;  they  might  be  termed  electro- 
neutral  elements.  It  should,  however,  be  stated  that  though  there 
is  a  distinct  parallelism,  there  is  no  direct  proportionality  between 
the  strength  of  a  base  and  the  affinity  of  the  metal  for  positive  elec- 
tricity (cf.  p.  435).  The  above  considerations,  which  are  of  funda- 
mental importance  for  the  proper  understanding  of  our  subject,  will 
now  be  illustrated  by  brief  references  to  the  main  groups  of  metals, 
arranged  on  the  basis  of  the  periodic  system. 

The  Principal  Groups  of  Metals— (i)  The  Alkali  group. 
The  closely  allied  elements,  lithium,  sodium,  potassium,  rubidium 
and  caesium,  which  occur  in  the  first  group  of  the  periodic  table, 
are  termed  the  alkali  metals.  They  are  strongly  electro-positive,  and 
this  character  increases  with  increasing  atomic  weight.  Correspond- 
ing with  this  they  are  powerful  bases  and  their  salts  are  not  hydro- 
lyzed in  solution.  They  act  exclusively  as  univalent  elements. 

(2)  The  Copper  group,  comprising  copper,  silver,  and  gold  also 
occur  in  the  first  vertical  column  of  the  periodic  table.  Corresponding 


GENERAL   PROPERTIES   OF   THE   METALS      389 

with  their  position,  they  can  all  function  as  univalent  elements,  but 
copper  and  gold  also  show  other  valencies.  The  hydroxides  of 
copper  and  gold  are  relatively  weak  bases,  and  their  salts  are  therefore 
partially  hydrolyzed  in  solution.  Silver  hydroxide  is  a  comparatively 
strong  base.  Their  affinity  for  positive  electricity  is  small. 

(3)  Metals   of  the    Alkaline    Earths  —  These    comprise    calcium, 
strontium,  and  barium.     They  always  act  as  bivalent  elements  ;  the 
hydroxides  are   strong   bases.      They  are   strongly   electro-positive 
elements. 

(4)  The  Zinc  group.     This  group  comprises  beryllium,  magnesium, 
zinc,  cadmium,  mercury.     Their  main  valency  is  2.     They  are  weaker 
bases  than  the  metals  of  the  alkaline  earths,  and  their  salts  with 
strong  acids  are  partially  hydrolyzed  in  solution. 

(5)  The   Aluminium   group.     Aluminium  is   the   only   important 
member   ot   this    family.      In    its    compounds   it   is    trivalent;    the 
hydroxide  has  weak  basic  and  also  weak  acidic  properties.     Never- 
theless, aluminium  is  a  strongly  electro-positive  element  (p.  481). 

(6)  The  Tin  group.     Tin  and  lead  are  the  most  important  members 
of  this  group.     Corresponding  with  their  position  in  the  table,  they 
are  quadrivalent,  but  both  can  also  function  as  bivalent  elements.     In 
the   divalent   condition  they  are  weakly  basic  ;  in  the  quadrivalent 
condition  the  acidic  character  predominates. 

(7)  The  Arsenic  group.     Arsenic,  antimony  and  bismuth  are  the 
chief  metals  belonging  to  this  group.     They  function  mainly  as  triva- 
lent and  pentavalent   elements.     The  metallic  character  increases 
markedly  with  increase  of  atomic  weight ;  but  even  bismuth  oxide  is 
a  weak  base,  and  its  salts  with  strong  acids  are  partially  hydrolyzed. 

(8)  The  Chromium  group.     Chromium,  the  chief  metal  in  the  sixth 
group,  functions  as  a  di-,  tri-,  quadri-  and  hexavalent  element.     The 
lowest  oxide  is  weakly  basic,  the  highest  acidic. 

(9)  The  Manganese  group.     Manganese  itself  is  the  only  metal  in 
this  group.     It  shows  a  number  of  valencies  from  2  to  7.     In  its  lowest 
state  of  valency  it  acts  as  a  weak  base,  in  its  highest  valency  it  has 
pronounced  acidic  properties. 

(10)  The  Iron  group.     Iron,  cobalt  and  nickel  are  the  members  of 
this  group.     They  act  as  divalent  and  trivalent  elements,  and  are 
weak  bases. 

(11)  The  Palladium  and  Platinum  groups.    The  members  of  these 
groups  show  considerable  diversity  in  valency.     The  lower  oxides  are 
weakly  basic,  the  higher  show  slight  acidic  properties 

From  the  above  brief  statement  of  the  characteristics  of  the  groups  , 


390    A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

certain  important  generalizations  may  be  made.  The  electro-positive 
character  (and  also  the  electro-negative  character)  is  most  pronounced 
for  univalent  elements  (e.g.  the  alkali  metals,  the  halogens) ;  but 
polyvalent  elements  may  also  show  great  affinity  for  electricity,  e.g. 
aluminium.  On  the  other  hand,  univalent  elements  are  mostly  strongly 
basic  or  acidic,  and  the  oxides  of  polyvalent  elements  are  often  both 
weakly  basic  and  weakly  acidic  (e.g.  aluminium).  When  the  same 
element  forms  both  basic  and  acidic  oxides,  the  higher  oxides  (that 
is,  those  containing  the  higher  proportion  of  oxygen)  are  most  acidic. 

Methods  of  Preparing:  Metals  from  their  Compounds 
— The  methods  will  be  described  in  connexion  with  the  metals  them- 
selves, but  a  brief  outline  of  the  processes  will  prove  useful  at  the 
present  stage.  The  compounds  chiefly  employed  are  the  oxides 
and  the  halides,  and  the  method  used  in  any  particular  case 
depends  mainly  on  the  volatility  of  the  metal  and  on  its  chemical 
affinity  for  other  elements.  Provided  the  oxide  is  not  too  stable,  it 
can  be  reduced  by  heating  with  carbon  (examples :  zinc,  tin,  iron,  etc.) 
or  with  finely  divided  aluminium  (examples  :  chromium,  manganese, 
etc.).  If  the  metal  has  so  great  an  affinity  for  oxygen  that  this 
method  cannot  be  applied,  recourse  is  often  had  to  the  electrolysis  of 
a  fused  salt,  for  example  the  fused  chloride,  the  metal  separating  at 
the  negative  pole.  Potassium,  sodium,  lithium,  magnesium,  alumi- 
nium and  other  metals  are  now  obtained  commercially  in  this  way. 

In  applying  these  general  methods  the  naturally  occurring  com- 
pound of  the  metal  has  first  to  be  converted  into  the  oxide  or  chloride, 
and  purification  may  also  be  necessary.  The  methods  used  depend 
largely  on  the  nature  of  the  ore.  Sulphides,  for  example,  are  con- 
verted into  oxides  by  roasting,  the  sulphur  being  burned  away  as 
sulphur  dioxide. 

Methods  of  Preparing:  Salts— The  more  important  methods 
of  preparing  salts  have  already  been  described  in  connexion  with  the 
individual  acids.  If  one  or  more  of  the  reacting  substances  are  used 
in  solution,  the  method  to  be  employed  in  any  particular  case  depends 
greatly  on  whether  the  salt  is  or  is  not  insoluble  in  water.  If  it  is 
only  slightly  soluble  the  general  method  of  preparation  is  by  double 
decomposition  between  a  solution  containing  the  basic  and  one 
containing  the  acidic  constituent,  e.g. : 

BaCl2+H2S04->BaS04  j  +2HC1 
Pb(NO3)2+2KI->PbI2i  +  2KNO3. 

If,  on  the  other  hand,  the  salt  is  soluble  in  water,  it  may  usually  be 


GENERAL   PROPERTIES   OF   THE    METALS      391 

prepared  by  the  action  of  the  corresponding  acid  (if  a  moderately 
strong  one)  on  the  metal,  oxide,  hydroxide,  or  salt  with  a  weak  or 
readily  volatile  acid.  Examples  : 


Ba(OH)2  +  2HCl-»BaCl2  +  2H 


2NaCl4-H2SO4->Na2SO4 

Some  salts  (e.g.  the  halides  or  sulphides)  are  occasionally  prepared 
by  direct  combination  of  the  elements. 

It  is,  of  course,  evident  that  salts  of  a  very  weak  base  and  a  very 
weak  acid  cannot  exist  in  contact  with  water,  as  they  readily  undergo 
hydrolysis.  Such  salts  can,  however,  sometimes  be  prepared  by 
reactions  in  the  dry  state. 

Chlorides  may  be  obtained  by  direct  combination  of  the  elements 
(e.g.  FeCl3  ;  SnCl4)  or  by  the  action  of  hydrochloric  acid  on  the  metal, 
oxide  or  carbonate.  The  three  chlorides  which  are  practically  in- 
soluble in  water  (silver,  lead  and  mercurous  chlorides)  are  prepared  by 
double  decomposition.  The  same  remarks  apply  to  bromides  and 
iodides. 

Carbonates,  being  all  insoluble  in  water  with  the  exception  of  those 
of  the  alkali  metals,  are  prepared  by  double  decomposition.  The 
alkali  carbonates  are  prepared  by  special  methods  (q.v.). 

Nitrates,  being  all  soluble  in  water,  are  prepared  by  the  methods 
used  for  obtaining  soluble  salts. 

Sulphates,  which  are  mostly  soluble  in  water,  are  prepared  by 
the  general  methods.  The  sulphates  of  calcium,  strontium,  barium, 
and  lead,  which  are  only  very  slightly  soluble  in  water,  are  prepared 
by  double  decomposition. 

Sulphides  —  Most  sulphides  are  insoluble  in  water,  and  are  pre- 
pared by  double  decomposition.  The  sulphides  of  the  alkalis,  which 
are  soluble,  are  obtained  by  the  action  of  hydrogen  sulphide  on 
the  corresponding  hydroxides. 

Phosphates,  with  the  exception  of  those  of  the  alkali  metals,  are 
insoluble  in  water,  and  are  obtained  by  double  decomposition. 
The  preparation  of  the  phosphates  of  the  alkalis  has  already  been 
described. 

Solubility  of  Salts  —  The  importance  of  a  knowledge  of  tne 
solubility  of  salts  in  water  has  already  been  repeatedly  emphasized. 
The  accompanying  table,  a  slightly  modified  form  of  one  given 


392     A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

by  Kohlrausch,  gives  the  solubility,  in  grams  per  litre  at  18°,  of 
most  of  the  bases  and  salts  in  general  use.  The  numbers  in  the 
squares  represent  the  solubility  of  the  salt  whose  positive  ion 
stands  at  the  top,  and  whose  negative  ion  stands  at  the  side  of 
the  column.  Each  square  contains  two  numbers  ;  the  top  one 
represents  the  number  of  grams  of  salt  (calculated  as  anhydrous) 
taken  up  by  1000  grams  (i  litre)  of  solvent  at  18°,  while  the  lower 
one  gives  the  number  of  grams  of  salt  in  i  litre  of  the  saturated 
solution  at  18°. 


£ 


o  o 

rf  xo    ^O  VO 


00  00      tx  tx 

do"    do* 


INO   ^  K 


d  d    o"  d 


O  O  d    CM  CO  CO 

O  O  CO  O    O 

o  q  q  q  o  q 

d  d  do  do 


88 

d  d 


d  d    d  d 


TOOO     t^t? 
to^-    O  oo 


0   0 

d  o 


0   0 

d  d 


d  d 


£vS      ?* 


2$  od 

CM    H      d    O 


tX     M          (JO 


"010      tx   tX       MM 
d*    CM        MM        d    d 


00  00 

to  10  co  co 

88  S5 

do  do 


2      CM    d 

d  d 


tx  tx 
tx  tx 


00         MM 


VO   VO          MM 

O   O      O  O 

do    do 


JaH 

CO  CO     H 


g|^«   t      S3- 

O    rt-     „    H      txvO       CO  CO 


CO  CO 

CM    d       CO  CO 

88  88 

c  d    d  d 


rt-  •<*•    co  co 

Cx  tx     CM    d 

d  d    d  d 


tO  CO     rj-  Tf 

do    do 


00  00 
0  0 
0  0 


ON 
%c<0 


O  O 

coco 


ON  CTi 

toto 

d  d 


8S8  £00  $$ 

^  co    tx  f    ^ 


ONOO      10 

M"  M"     ^S) 


do    do    do 


z&  3" 

S  2"  2 


§§ 


%"cO      g^g5 

dd    oo' 


II M 


rr  O       2» 
^&     CO 


^a  -  H 


00    Cx     Cx  O 

tO    M       00      M 

CO  CO    00    Cx 


•*    d  to 

VO       M    CO 
M     VO    tO 


d  O 

i  CO     O\  O^ 
I  CO      M    M 


£&£ 

CO  M     VO 


as. 


^   5§    o| 


o^<§ 

CO  d       H 


M  M 


cc  _w 

o     o 

U        ffl 


•&      O    '  J* 

O       tvo       U 


o5 

U 


GENERAL   PROPERTIES   OF   THE   METALS      393 

The  table  just  given  requires  to  be  amplified  by  a  statement  as 
to  whether  the  compound  in  equilibrium  with  the  saturated  solution 
at  1 8°  is  the  anhydrous  salt  or  a  hydrate.  This  information  is  given 
in  the  accompanying  table,  also  clue  to  Kohlrausch.  The  integers 
give  the  number  of  molecules  of  water  associated  with  the  salt 
whose  positive  component  is  represented  at  the  top  and  negative 
component  at  the  side  of  the  diagram.  Thus  barium  chloride  at 
1 8°  separates  from  solution  with  2H2O.  The  letters  give  further 
information  as  to  the  behaviour  of  the  salt ;  h  denotes  that  it  is 
stable  on  heating,  a  that  it  is  stable  in  air,  e  that  it  is  efflorescent,  and 
d  that  it  is  deliquescent  (p.  402). 


K 

Na 

Li 

Ag 

Ba 

Sr 

Ca 

Mg 

Zn 

Cd 

Cu 

Pb 

Cl    .     . 

oh 

oh 

oh 

oh 

2  a 

6e 

6d 

6d 

ii* 

4« 

2d 

oh 

Br   .     . 

oh 

oh 

oh 

ok 

2  a 

6d 

6d 

6d 

2d 

4* 

4* 

oh 

I      .     . 

Oh 

oh 

oh 

oh 

2d 

6d 

od 

Sd 

od 

o  a 

oh 

NOS     . 

o  a 

o  a 

o  a 

o  a 

o  a 

4* 

*d 

6d 

6d 

4d 

6d 

o  a 

CIO.,    . 

o  a 

o  a 

od 

o  a 

i  a 

Sd 

2d 

6d 

6d 

2d 

6d 

i  a 

Br03.   . 

o  a 

o  a 

od 

o  a 

i  a 

i  a 

i  a 

6e 

6a 

2  a 

6a 

i  a 

10,.     . 

o  a 

o  a 

od 

o  a 

i  a 

6e 

6e 

4« 

2  a 

o  a 

i  a 

o  a 

C2H302 

od 

3d 

2d 

o  a 

i  a 

\« 

2d 

4d 

$a. 

3d 

i  a 

3<* 

S04.    . 

oh 

10  e 

oh 

o  a 

oh 

oh 

2  a 

7  e 

7  e 

Q/3a 

•s« 

oh 

Cr04    . 

oh 

10  e 

2  a 

o  a 

oh 

oh 

i  a 

7  e 

oh 

C204    . 

i  a 

oa 

o  a 

o  a 

i  a 

o  a 

i  a 

2  a 

2  a 

3" 

i  a 

o  a 

C03     . 

** 

10  e 

o  a 

o  e 

oh 

oh 

o  a 

3* 

o  a 

o  a 

o  a 

CHAPTER   XXVI 
ELEMENTS   OF   GROUP    I.,   SUB-GROUP  A 


sub-group   includes  the  following  five  metals,  which  are 
J-      known  as  the  alkali  metals  :  — 

Atomic  Weight 
Lithium  (Li)         .....        •        6.94 

Sodium  (Na)         ......      23.00 

Potassium  (K)      ......      39-  IO 

Rubidium  (Rb)     ......      85.45 

Caesium  (Cs)       ......     132.81 

They  constitute  a  typical  family  of  elements,  inasmuch  as  they 
resemble  each  other  very  closely  in  their  physical  and  chemical 
properties.  As  in  the  case  of  other  families  of  elements,  these 
properties  vary  regularly  with  increasing  atomic  weight.  They  are 
all  soft  metals,  which  attack  water  at  the  ordinary  temperature,  and 
have  very  great  affinity  for  oxygen.  They  act  as  strongly  electro- 
positive univalent  elements,  and  the  electro-positive  character  in- 
creases with  increase  in  atomic  weight.  Nearly  all  the  salts  of  the 
alkalis  are  readily  soluble  in  water.  The  salts  of  the  different  metals 
arts  distinguished  by  the  characteristic  colours  imparted  to  a  flame 
(cf.  p.  408).  The  most  important  members  of  the  group  are  sodium 
and  potassium  ;  rubidium  and  caesium  are  very  rare  elements.  The 
general  properties  of  the  members  of  this  sub-group  are  compared  in 
detail  at  the  end  of  the  chaper. 

SODIUM 

Symbol,  Na.     Atomic  weight=23.     Molecular  weight=23  (doubtful). 

History  —  Of  the  compounds  of  sodium  the  carbonate  has  been 
known  for  many  centuries,  and  is  probably  referred  to  in  the  Bible  under 
the  name  nitre.  In  the  sixteenth  century  the  Arabs  introduced  into 
Europe  the  name  "  natron  "  for  sodium  carbonate  in  contrast  to  nitre, 
which  was  applied  to  potassium  nitrate.  At  that  time,  however,  no 

394 


SODIUM 


395 


distinction  was  drawn  between  sodium  salts  and  potassium  salts. 
The  latter  were  chiefly  obtained  from  vegetable  sources,  and  were 
called  kali  or  alkali.  This  distinction  was  first  drawn  by  Duhamel 
du  Monceau  in  1736. 

Occurrence — Sodium  does  not  occur  free  in  nature  owing  to  its 
great  affinity  for  oxygen.  It  occurs  as  silicate  in  many  rocks,  by  the 
disintegration  of  which  it  finds  its  way  into  the  soil  and  thence  into 
rivers  and  the  sea.  Sodium  chloride  occurs  to  the  extent  of  2.6  to 
2.9  per  cent,  in  sea-water,  and  the  same  salt  is  found  in  deposits  of 


FIG.  81 

great  thickness  in  various  parts  of  the  world  ;  for  example,  in  Russia, 
in  Germany,  and  in  Lancashire  and  Cheshire  in  this  country. 
Sodium  nitrate  occurs  in  enormous  amount  in  Chili  and  Peru,  and  is 
termed  Chili  saltpetre.  Large  deposits  of  cryolite,  sodium  aluminium 
fluoride,  AlF3,3NaF,  occur  in  Greenland  and  Iceland, 

Preparation — Metallic  sodium  was  first  obtained  by  Davy 
(1807)  by  the  electrolysis  of  sodium  hydroxide.  It  is  an  interesting 
fact  that  at  present  the  metal  is  prepared  on  the  commercial  scale 
almost  entirely  by  this  method. 

(i)  Cosiness  Electrolytic  Process— The  method  for  this  purpose 
devised  by  Castner  is  illustrated  in  Fig.  81.  D  is  a  cylindrical  steel 


396     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

vessel  with  an  opening  at  the  bottom,  through  which  the  iron  cathode 
C  passes  ;  it  is  heated  by  gas-burners  e  e,  so  that  all  except  the  neck 
is  kept  at  a  temperature  about  20°  above  the  melting-point  of  the 
hydroxide.  The  anode  A,  which  is  conveniently  made  in  the  form  of 
a  cylinder  with  vertical  slits,  surrounds  the  upper  portion  of  the 
cathode.  Within  the  anode  is  the  collecting  pot  F,  from  which  is 
suspended  a  cylinder  of  wire  gauze,  which  surrounds  the  upper 
portion  of  the  cathode. 

The  products  of  electrolysis  are  sodium  and  hydrogen,  which  are 
liberated  at  the  cathode,  and  oxygen,  which  is  liberated  at  the  anode, 
and  escapes  through  a  valve  at  the  top  of  the  vessel.  The  sodium 
rises  to  the  surface  of  the  fused  electrolyte  in  F,  and  is  removed  from 
time  to  time  by  means  of  a  perforated  ladle.  The  hydrogen,  also 
liberated  at  the  cathode,  serves  to  protect  the  sodium  against  oxida- 
tion ;  it  escapes  at  the  loosely-fitting  lid  of  the  vessel  F. 

(2)  Other  Electrolytic  Methods  —  Instead  of  using  sodium  hydroxide, 
it  would  obviously  be  advantageous  in  some  ways  if  sodium  chloride 
could  be  used  as  electrolyte  :    it  is  cheaper,  and  both  products  of 
electrolysis,  sodium  and  chlorine,  are  commercially  valuable.    Among 
the  difficulties  met  with  in  using  this  method  are  the  high  melting- 
point  of  the  chloride,  the  disintegrating  effect  of  the  fused  electrolyte 
on  the  cell  materials,  and  the  combination  of  the  liberated  metal  with 
the  fused  salt  to  form  the  "  sub-chloride,"  Na2Cl.     These  difficulties 
do  not  appear  to  have  been  entirely  overcome. 

(3)  Chemical   Methods  —  The    chemical    methods    for    preparing 
metallic  sodium,  which  have  now  been  almost  completely  displaced 
by  the  electrolytic  methods,  depend  upon  the  reduction  of  sodium 
hydroxide  or  carbonate  with  carbon  or  iron,  or  a  mixture  of  both.     In 
Castner's  chemical  process  sodium  hydroxide  was  mixed  with  iron 
and  finely  divided   carbon   (perhaps   iron   carbide,   FeC2),  and  the 
mixture  distilled  from  iron  retorts  : 


The  metal  was  condensed  in  flat  iron  receivers,  with  an  outflow  so 
arranged  that  the  metal  was  collected  under  mineral  oil. 

Physical  Properties  —  Sodium  is  a  white,  lustrous  metal, 
which  at  room  temperature  can  be  moulded  with  the  fingers,  but  is  hard 
at  -20°.  It  melts  at  97°,  and  boils  at  877°;  the  vapour  in  thick 
layers  has  a  purplish  colour.  The  vapour  density,  determined  in 
platinum  vessels,  is  about  12,  indicating  that  the  metal  is  monatomic 
m  this  state. 


SODIUM  397 

Chemical  Properties  —  Sodium  is  unattacked  in  perfectly  dry 
air  or  oxygen,  but  in  moist  air  the  fresh  surface  becomes  coated 
almost  instantaneously  with  a  film  of  oxide.  It  is  vigorously  acted 
upon  by  water  at  room  temperature  : 


The  heat  given  out  in  the  reaction  is  not  sufficient  to  ignite  the 
hydrogen  if  the  metal  is  allowed  to  move  about  on  the  surface  of  the 
water,  but  if  the  metal  is  confined  'to  one  point  the  hydrogen  ignites 
and  burns  with  a  yellow  flame. 

Sodium  dissolves  in  liquid  ammonia  to  form  a  blue  liquid.  It  com- 
bines with  dry  ammonia  at  300-400°  to  form  sodamide,  NaNH2,  a 
white  solid  which  melts  at  155°.  With  mercury  in  certain  proportions 
it  forms  a  solid  amalgam  (p.  472),  which  contains  definite  compounds 
of  the  two  elements.  The  action  of  water  on  sodium  amalgam  forms 
a  convenient  method  of  obtaining  "  nascent  "  hydrogen.  With  potas- 
sium, sodium  forms  a  liquid  alloy  resembling  mercury,  which  has 
been  used  for  thermometric  purposes. 

Sodium  Hydride,  NaH  —  When  sodium  is  heated  with  hydrogen 
in  an  iron  vessel  at  360°,  sodium  hydride  is  formed,  and  condenses  on 
a  cooler  part  of  the  apparatus  in  the  form  of  colourless  crystals.  The 
compound  readily  dissociates  on  heating  above  430°.  It  is  stable  in 
dry  air,  but  is  immediately  decomposed  by  water. 

Sodium  Oxides  —  Two  oxides  are  known,  the  normal  oxide 
or  monoxide,  Na2O,  and  the  peroxide,  Na2O2. 

Sodium  Monoxide  is  obtained  by  partial  oxidation  of  sodium, 
the  unchanged  metal  being  removed  by  prolonged  distillation  in 
a  vacuum.  It  is  a  white,  amorphous,  hygroscopic  powder,  which 
combines  vigorously  with  water  to  form  the  hydroxide. 

Sodium  Peroxide,  Na2O2,  is  obtained  by  heating  sodium  in 
an  iron  tube  at  300°  in  a  current  of  dry  air,  free  from  carbon  dioxide. 
It  occurs  as  a  light  yellow  powder,  which  dissolves  in  water  under 
ordinary  conditions  with  rise  of  temperature  and  considerable  evolu- 
tion of  oxygen.  The  primary  reaction  is  probably  as  follows  :  — 


but  unless  precautions  are  taken  to  keep  the  mixture  cool,  the 
hydrogen  peroxide  is  largely  decomposed  into  water  and  oxygen.  The 
above  reaction  is  reversible.  When  the  peroxide  is  dissolved  in 
dilute  acids,  the  corresponding  sodium  salt  and  hydrogen  peroxide 
are  obtained. 


398     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

An  octahydrate,  Na2O2,8H2O,  has  been  obtained  in  colourless 
crystals.  Sodium  peroxide  is  largely  used  for  oxidizing  and  bleach- 
ing purposes. 

Sodium  Hydroxide  (Caustic  Soda),  NaOH—  Preparation—  (i) 
By  the  action  of  water  on  metallic  sodium  or  sodium  monoxide.  The 
former  method  is  used  for  the  preparation  of  very  pure  hydroxide. 

(2)  By  boiling  together  sodium  carbonate  and  calcium  hydroxide 
in  iron  vessels  : 


The  reaction  is  a  reversible  one,  and  in  practice  the  best  yield  is 
secured  by  using  excess  of  calcium  hydroxide  and  working  at  a 
definite  dilution.  The  calcium  carbonate  is  allowed  to  settle,  the 
clear  liquid  decanted  off,  evaporated  in  iron  vessels,  and  cast  into 
sticks. 

(3)  Caustic  soda  is  now  obtained  commercially  by  electrolysis  of 
sodium  chloride.  The  methods  differ  according  as  the  salt  is  em- 
ployed in  the  fused  condition  or  in  aqueous  solution.  As  an  illustra- 
tion of  the  former  case  the  Acker  process,  formerly  in  use  at  Niagara 
Falls,  may  be  referred  to.  The  salt  is  fused  by  the  heat  of  the  cur- 
rent itself,  and  is  electrolyzed  between  a  carbon  anode  and  a  lead 
cathode.  The  alloy  of  sodium  ^nd  lead  thus  obtained  is  caused  to 
flow  into  another  compartment,  where  it  is  decomposed  by  a  jet  of 
steam  which  forces  the  sodium  hydroxide  and  lead  into  a  third  com- 
partment, from  which  the  latter  is  returned  to  the  cell  and  the  former 
run  off. 

The  methods  depending  upon  the  electrolysis  of  sodium  chloride 
in  solution  are  mainly  of  two  types  :  (a)  the  sodium  is  liberated  at  a 
mercury  cathode,  the  amalgam  removed  from  the  cell  and  decom- 
posed in  a  separate  compartment  ;  (b)  the  sodium  is  converted  into 
alkali  in  the  cell  itself,  a  diaphragm  being  used  to  separate  the  anode 
from  the  cathode  department.  Space  will  only  admit  of  a  brief  ac- 
count of  the  Solvay  mercury  cell,  used  both  in  Belgium  and  America. 
A  diagram  of  the  apparatus  is  given  in  Fig.  82.  It  consists  of  a  large, 
slightly  tilted  rectangular  cement  trough,  through  which  streams  of 
mercury  and  brine  flow  continuously.  The  brine  enters  at  Sx  and 
flows  out  at  S2,  the  mercury  enters  at  Mt  and  the  amalgam  flows  out 
at  M2.  The  anodes  are  of  carbon  (or  platinum),  and  are  placed  1-2 
cm.  from  the  upper  surface  of  the  mercury  cathode.  The  issuing 
amalgam  is  decomposed  by  water  in  a  separate  vessel  and  the  mer- 
cury returned  to  the  cell, 


SODIUM 


399 


Properties  —  Sodium  hydroxide  is  a  white,  very  hygroscopic 
substance,  which  melts  at  318°.  It  liquefies  in  the  air  owing  to 
absorption  of  moisture,  but  finally  becomes  solid  consequent  on  for- 

.  I          .  /^.  -»  CHLORINE 


£  K 


FIG.  82. 


mation  of  the  carbonate.     Its  aqueous  solution  has  powerful  caustic 
properties,  and  shows  all  the  characteristics  of  a  strong  base. 

The  discovery  of  the  true  relationship  between  alkali  (and  alkaline  earthj 
carbonates  and  hydroxides  is  due  to  Joseph  Black  (cf.  p.  335).  Before  his  time 
the  former  were  regarded  as  simple  substances,  and  it  was  assumed  that  when, 
e.g.,  limestone  was  calcined  firestuff  (phlogiston)  was  taken  up,  and  when  the 
alkali  carbonates  were  causticised  by  boiling  with  lime  the  phlogiston  was  trans- 
ferred to  the  former.  Black  showed  that  when  limestone  or  magnesium  carbonate 
is  heated,  a  gas  which  he  termed  fixed  air,  is  given  off  with  consequent  loss  of 
weight,  that  this  gas  could  be  fixed  by  the  caustic  alkalis  (forming  carbonates), 
and  that  the  resulting  compounds  again  became  caustic  when  their  carbon  dioxide 
was  removed  by  lime  or  magnesia. 

Sodium  Chloride,  NaCl — As  already  mentioned,  sodium  chlo- 
ride is  the  chief  saline  constituent  of  sea-water,  and  it  occurs  in  large 
proportions  in  certain  lakes  and  also  in  deposits  of  great  thickness  in 
Galicia,  Germany,  and  other  parts  of  the  world.  These  deposits 
were  formed  by  the  complete  drying  up  of  lakes  or  seas.  As  sodium 
chloride  is  less  soluble  than  many  of  the  other  salts  in  sea-water,  it 
separates  out  first,  and  above  it  are  found  layers  of  more  soluble  salts. 
The  most  famous  of  such  deposits  is  at  Stassfurt  in  Germany,  where 
the  salt  layers  are  in  some  parts  1000  metres  in  thickness. 

Salt  is  prepared  in  warm  countries  from  sea-water  by  evaporation, 
and  from  salt  deposits  either  by  mining  the  solid  or  by  boring  through 
the  upper  strata  and  then  using  water  to  form  a  strong  brine,  which 
is  pumped  through  copper  tubes  and  evaporated  to  crystallization. 
In  order  to  economize  fuel,  the  first  stage  of  the  evaporation  is  often 
accomplished  by  allowing  the  brine  to  trickle  down  large  ricks  of 
brushwood,  so  that  free  exposure  to  wind  is  secured.  In  order  to 


400    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

render  it  fit  for  household  purposes,  natural  salt  has  to  be  purified  by 
recrystallization  from  water.  The  deliquescent  character  of  some 
samples  of  salt  is  due  to  the  presence  of  magnesium  chloride  as 
impurity.  Sodium  chloride  can  be  obtained  in  a  perfectly  pure 
condition  by  precipitating  from  aqueous  solution  by  means  of  hydrogen 
chloride  (cf.  p.  439)  and  then  heating  strongly. 

Properties — Sodium  chloride  is  usually  obtained  in  cubic  crystals, 
which  melt  about  810°.  The  solubility  in  grams  per  100  grams  of 
water  is  as  follows  :  35.63  at  o°,  35.82  at  20°,  36.32  at  40°,  37.06  at  60°, 
and  39.12  at  100°.  The  very  slight  change  of  solubility  with  tempera- 
ture is  in  accordance  with  the  very  small  heat  of  solution. 

Only  one  hydrate  of  sodium  chloride,  NaCl,2H2O,  is  known  ; 
it  separates  from  solutions  below  o°  in  monoclinic  crystals,  whilst  at 
higher  temperatures  the  anhydrous  salt  is  obtained  in  cubic  crystals. 
The  fact  that  sodium  chloride  is  an  essential  constituent  of  the  food 
of  animals  is  familiar  to  all.  This  salt  is  the  source  of  all  sodium 
compounds  except  the  nitrate,  and  is  the  chief  source  of  chlorine 
compounds. 

Sodium  Bromide,  NaBr,  and  Sodium  Iodide,  Nal,  are 
prepared  by  the  methods  described  under  the  corresponding  potas- 
sium salts.  They  are  isomorphous  with  sodium  chloride.  At  room 
temperature  the  dihydrates  separate  from  aqueous  solution.  The 
transition  temperature,  NaBr-NaBr,2H2O,  is  at  50.6° ;  that  of 
NaI-NaI,2H2O  at  65°.  At  low  temperatures  sodium  iodide  pen- 
tahydrate  separates  from  solution ;  the  transition  temperature, 
NaI,5H2O-NaI,2H2O,  is  at  -  13.5°. 

Composition  of  Hydrates— The  general  methods  used  in 
determining  the  composition  of  the  hydrates  formed  by  a  salt  or  other 
substance  and  the  limits  within  which  they  exist  may  conveniently  be 
considered  in  connexion  with  sodium  iodide.  The  diagram  represent- 
ing the  variation  in  the  solubility  of  the  salt  with  temperature  is 
shown  in  Fig.  83  ;  the  ordinates  represent  temperatures  and  the 
abscissae  concentrations  in  grams  of  salt  per  100  grams  of  water.  The 
curve  AB  represents  the  effect  of  the  salt  in  lowering  the  freezing- 
point  of  water ;  in  other  words,  it  is  the  curve  along  which  ice  and 
solution  are  in  equilibrium  (cf.  the  system,  water-potassium  iodide 
(p.  199)).  When  sufficient  salt  is  added,  however,  a  point  is  reached 
at  which,  on  cooling  the  solution  sufficiently,  a  hydrate,  and  not  ice, 
separates.  Thus,  along  the  part  BC  of  the  curve  the  pentahydrate 
separates  from  solution  ;  in  other  words,  BC  represents  the  solubility 
curve  of  the  compound  NaI,5H2O.  At  the  point  B,  -  32.5°,  ice  and  the 


SODIUM 


401 


pentahyclrate  are  in  equilibrium,  and  this  is  a  eutectic  point.  As  the 
quantity  of  salt  is  still  further  increased  the  point  C  is  reached,  at 
-13.5°,  when  the  pentahydrate  and  dihydrate  are  in  equilibrium. 
With  higher  proportions  of  salt  the  solid  in  equilibrium  with  the  solu- 
tion is  the  dihydrate,  and  its  solubility  curve  is  represented  by  CD. 
Finally,  at  D,  65°,  the  dihydrate  is  in  equilibrium  with  anhydrous 
salt,  and  the  further  part  of  the  curve  DE  represents  the  solubility  of 


3*1 

1 

/ 

V 

50 

g.o 

r 

1- 

10 
0 

—  to 
\ 

-30 

"40 
< 

/ 

V 

*y 

\ 

/ 

*X 

X' 

^. 

1 

\ 

\! 

^ 

c 

\ 

V 

5 

SO          100          I&O         ZOO         250         300         3bO 

Concentration  (grams  in  100  grams  water). 
FIG.  83. 

the  anhydrous  salt.  At  each  of  the  points,  B,  C,  and  D,  so-called 
breaks  on  the  curve,  a  new  phase  appears  ;  and  as  the  sections  BC, 
CD,  DE  each  represent  the  solubility  curve  of  a  separate  substance, 
it  is  easy  to  understand  why  the  whole  curve  is  discontinuous.  The 
number  and  composition  of  the  hydrates  can  then  be  determined 
from  the  number  and  position  of  the  breaks  on  the  curve,  controlled 
by  analysis  of  the  compounds  in  equilibrium  with  the  solution  at 
different  dilutions. 
26 


402     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Formerly  the  number  and  nature  of  the  hydrates  formed  by  a  salt  was  deter- 
mined by  evaporating  the  solution  at  different  temperatures,  and  analysing  the 
solid  compounds  obtained.  The  modern  method  is  to  start  with  pure  water, 
add  successive  portions  of  the  salt,  and  determine  the  complete  solubility  curve, 
as  shown  in  Fig.  83.  If  no  hydrates  exist  under  the  experimental  conditions, 
the  curve  representing  the  equilibrium  between  salt  and  solution  will  be  con- 
tinuous, as  in  the  examples  in  Fig.  32.  The  existence  of  hydrates  will  be 
indicated  by  breaks  in  the  curve.  The  complete  curve  of  the  system  includes 
the  eutectic  ice-salt  and  that  part  of  the  curve  where  the  solution  is  in  equilibrium 
with  ice,  but  this  part  is  generally  omitted  in  the  solubility  curves. 

Other  properties  of  the  system,  besides  the  solubility,  may  also  be  made  use 
of  to  determine  the  existence  and  composition  of  hydrates.  One  of  the  most 
instructive  of  these  properties  is  the  vapour  pressure  of  solid  hydrates.  If  a 
crystal  of  copper  sulphate  pentahydrate  is  put  in  the  vacuum  of  the  barometric 
column  (Fig.  19)  it  will  be  found  to  exert  a  definite  pressure,  due  to  water 
vapour,  just  as  water  itself  does  under  the  same  conditions.  At  50°  the  pressure 
in  question  is  about  47  mm.  If  the  copper  sulphate  is  put  in  a  desiccator  over 
concentrated  sulphuric  acid,  which  absorbs  the  vapour,  it  will  be  found  that  the 
pressure  remains  constant  at  47  mm.  (if  the  observations  are  made  at  50°)  till  the 
salt  has  lost  2H2O,  when  it  suddenly  falls  to  30  mm. ,  which  is  the  vapour  pressure 
of  the  trihydrate.  As  long  as  any  trihydrate  is  present  it  remains  constant  at 
this  value,  but  when  two  further  molecules  of  water  have  been  removed  the 
pressure  suddenly  falls  to  4.4  mm.,  which  is  the  vapour  pressure  of  the  mono- 
hydrate.  Finally,  when  the  water  is  completely  removed,  the  pressure  of  course 
falls  to  zero. 

It  is  evident  from  the  above  that  the  occurrence  of  intermediate  hydrates  can 
be  deduced  from  observation  of  the  pressure  during  dehydration.  If  the  pressure 
fell  at  once  from  47  mm.  to  zero,  it  would  mean  that  no  intermediate  hydrates 
exist. 

Conversely,  if  water  vapour  at  a  pressure  of  say  5  mm.  is  brought  into  contact 
with  anhydrous  copper  sulphate  at  50°,  the  monohydrate  will  be  formed,  but  no 
higher  hydrate.  Only  when  the  pressure  is  raised  above  30  mm.  is  the  trihydrate 
formed,  and  a  pressure  exceeding  47  rnm.  is  required  to  form  the  pentahydrate. 
The  reactions  in  question  are  therefore  reversible,  and  the  phenomenon  is  one  of 
dissociation. 

These  considerations  enable  us  to  understand  the  behaviour  of  a  hydrated  salt 
in  the  air.  If  the  pressure  of  aqueous  vapour  exerted  by  the  hydrate  is  greater 
than  the  partial  pressure  of  aqueous  vapour  in  the  atmosphere,  the  salt  will  lose 
water  and  a  lower  hydrate  will  be  formed.  Under  these  circumstances  the  salt  is 
said  to  effloresce  in  the  air.  At  room  temperature,  the  vapour  pressure  of  copper 
sulphate  pentahydrate  is  less  than  the  average  pressure  of  aqueous  vapour  in  the 
atmosphere,  and  the  salt  does  not  effloresce  ;  the  vapour  pressure  of  sodium 
sulphate  decahydrate,  on  the  other  hand,  is  greater  than  the  average  pressure  of 
aqueous  vapour  in  the  atmosphere,  and  the  salt  effloresces. 

It  should  be  pointed  out  that  although  we  have  spoken  of  47  mm.  as  the  dis- 
sociation pressure  of  copper  sulphate  pentahydrate  at  50°,  a  definite  pressure  is 
only  attained  when  two  solid  phases,  the  hydrate  in  question  and  the  next  lower 
hydrate,  are  present.  Thus,  at  50°,  copper  sulphate  trihydrate  is  stable  in  the 
presence  of  pressures  of  aqueous  vapour  from  a  little  above  30  mm.  to  a  little 


SODIUM 


403 


below  47  mm.  ;  when  it  is  kept  at  a  value  less  than  30  mm.  the  trihydrate  finally 
disappears  and  monohydrate  is  formed  exclusively  ;  when  it  is  kept  at  a  value 
above  47  mm.  ,  the  pentahydrate  is  finally  obtained.  Many  hydrates  thus  remain 
unaltered  in  the  air  when  the  temperature  and  vapour  pressure  are  altered  within 
certain  limits. 

When  the  vapour  pressure  of  the  saturated  solution  of  a  salt  is  less  than  the 
vapour  pressure  of  the  atmosphere,  the  salt  takes  up  water,  and  a  solution  is 
finally  obtained.  Such  a  salt  is  said  to  be  deliquescent—  example,  calcium 
chloride. 

From  the  general  principles  of  equilibrium,  the  dissociation  pressure  of  a 
hydrated  salt  must  evidently  correspond  with  the  vapour  pressure  of  the  aqueous 
solution  with  which  it  is  in  equilibrium.  Thus  sodium  iodide  dihydrate  separates 
from  aqueous  solution  at  all  temperatures  between  -  13°  and  +67°,  and  we  may 
anticipate,  as  is  in  fact  the  case,  that  it  remains  unaltered  in  pressures  of  aqueous 
vapour  from  about  2  mm.  (the  pressure  at  —  13°)  to  200  mm.  (the  pressure  at  67°), 
and  therefore  that  it  is  stable  in  the  air. 

Sodium  Carbonate,  Na2CO3  —  Preparation  —  This  salt, 
which  is  extensively  used  in  the  manufacture  of  glass,  soap,  etc.,  is 
now  prepared  commercially  from  sodium  chloride  by  three  distinct 
methods  —  (i)  The  Leblanc  process,  invented  about  1790  ;  (2)  the 
ammonia-soda  or  Solvay  process  (1860)  ;  (3)  the  electrolytic  process, 
introduced  within  the  last  few  years. 

(i)  The  Leblanc  Process  —  There  are  three  distinct  stages  in  this 
process.  In  the  first,  sodium  chloride  is  converted  into  the  sulphate 
by  heating  with  sulphuric  acid  (salt-cake  process)  ;  the  sulphate  is 
then  reduced  to  the  sulphide  by  heating  with  carbon,  and  finally, 
from  the  sulphide  and  calcium  carbonate  (chalk)  at  a  high  tem- 
perature sodium  carbonate  and  calcium  sulphide  are  obtained  : 


(a) 
(b) 
(c)  Na2S  +  CaCO3->Na2CO3+CaS. 

The  last  two  stages  are  accomplished  in  one  operation,  the  sodium 
sulphate  being  mixed  with  chalk  and  small  coal,  and  the  mixture 
strongly  heated.  The  product  (black  ash)  is  then  treated  with  water, 
which  dissolves  out  the  carbonate,  leaving  the  insoluble  calcium 
sulphide. 

With  reference  to  the  details  of  the  operation,  the  first  stage  of  the 
salt-cake  process  is  carried  out  by  heating  the  sak  and  sulphuric 
acid  in  large  cast-iron  pans  (Fig.  84,  d).  The  main  reaction  at  this 
stage  is  as  follows  :  — 

NaCl-f-H2SO4->NaHSO4 


404     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

The  practically  dry  product,  containing  sodium  bisulphate  and  un- 
altered chloride,  is  raked  out  on  the  hearth  (b]  of  a  reverberatory 
furnace1  (Fig.  84)  and  heated  still  more  strongly,  with  continual 
mixing,  until  the  following  reaction  is  complete  : 

NaHSO4  +  NaCl->Na2SO4  +  HCL 

The  hydrogen  chloride  escaping  in  the  two  stages  of  the  process  is 
absorbed  by  passing  it  up  towers  filled  with  coke  over  which  water 
is  made  to  trickle. 

The  production  of  the  black  ash  is  brought  about  by  heating  the 
reacting  substances  in  a  rotating  cylinder  (to  secure  thorough  mixing) 
in  a  furnace,  and  the  end  of  the  reaction  is  recognized  by  the  appear- 


FIG.  84, 

ance  of  jets  of  burning  carbon  monoxide,  due  to  the  action  of  carbon 
on  the  calcium  carbonate : 

CaCO3+  C-»CaO  +  2CO. 

The  carbon  monoxide  formed  towards  the  end  of  the  operation  has 
the  effect  of  rendering  the  black  ash  porous,  and  therefore  more 
readily  extracted  by  water. 

1  A  reverberatory  furnace  is  an  arrangement  whereby  the  material  to  be  heated 
or  fused  does  not  come  into  contact  with  the  solid  fuel.  A  simple  furnace  of  this 
type  is  shown  in  Fig.  85.  The  fuel  is  burned  in  the  fire-box  a,  and  the  flame  and 
heated  gases  are  caused  to  act  on  the  material  spread  out  on  the  bed,  c,  of  the 
furnace  by  "reverberation"  from  the  low  roof.  The  products  of  combustion 
escape  by  the  flue  d.  The  furnace  represented  in  Fig.  84  has  in  addition  the 
large  cast-iron  pan  d,  in  which  the  mixture  of  sodium  chloride  and  sulphuric 
acid  is  heated. 


SODIUM 


405 


The  extraction  (lixiviation)  of  the  black  ash  is  carried  on  in  a 
series  of  vessels  so  arranged  that  the  water  passes  in  turn  from  one 
to  the  other  until  it  is  completely  saturated.  The  process  is  carried 
on  at  30°  to  40°,  the  temperature  at  which  sodium  carbonate  is  most 
soluble.  The  solution  is  evaporated  in  vessels  heated  in  the  flues  of 
the  black-ash  furnace,  when  the  monohydrate,  Na2CO3,H2O,  is  pre- 
cipitated. When  the  monohydrate  is  strongly  heated,  all  the  water 
is  driven  off,  and  "calcined  soda,"  largely  used  in  commerce,  is 
obtained.  On  dissolving  the  anhydrous  salt  in  water  and  allowing 
it  to  crystallize,  the  readily  soluble  decahydrate,  Na2CO3,ioH2O, 
"  washing  soda,"  is  obtained  in  large  crystals. 

The  insoluble  residue  from  the  black  ash,  which  consists  largely 
of  calcium  sulphide,  is  worked  up  in  order  to  extract  the  sulphur. 
According  to  Chance's  process  for  this  purpose,  the  residue  mixed 
with  water  is  placed  in  a  series  of  vessels  through  which  carbon 
dioxide  is  passed.  The  final  result  of  the  reactions  is  that  calcium 
carbonate  and  a  gas  rich  in  hydrogen  sulphide  are  obtained.  The 
latter  is  either  burned  completely  to  sulphur  dioxide,  which  is  used 
directly  in  the  sulphuric  acid  manufacture  (p.  304),  or  is  burned  in 
an  insufficient  supply  of  air  and  the  sulphur  collected  : 


(2)  The  Ammonia-Soda  Process  —  According  to  this  simple  and 
economical  process,  introduced  by  Solvay,  sodium  chloride  and 
ammonium  bicarbonate  are  brought  together  in  fairly  concentrated 
solution,  when  sodium  bicarbonate,  NuHCO3,  being  relatively  slightly 
soluble  in  water,  is  precipitated,  and  is  then  converted  into  the 
normal  carbonate  by  heating.  In  practice,  ammonia  and  carbon 
dioxide  (obtained  by  heating  limestone)  are  led  into  a  solution  of 
sodium  chloride.  The  reactions  are  represented  by  the  following 
equations  :  — 

(i)  CO2  +  NH3+H2O-»NH4HCO3. 

(2) 
(3)  2 


From  the  ammonium  chloride  formed  in  the  course  of  the  reaction 
the  ammonia  is  regenerated  and,  along  with  the  carbon  dioxide 
formed  in  (3),  is  used  to  decompose  a  fresh  quantity  of  sodium 
chloride,  so  that  the  only  by-product  is  calcium  chloride.  The  process 
is  thus  a  very  economical  one,  and  it  has  the  additional  advantage 


406    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

over  the  Leblanc  process  that  the  soda  obtained  is  much  purer.  It 
must  be  clearly  understood  that  the  whole  basis  of  the  process  is 
the  relative  insolubility  of  sodium  bicarbonate  as  compared  with 
ammonium  bicarbonate. 

(3)  Electrolytic  Method  —  Sodium  hydroxide,  obtained  by  an  electro- 
lytic method  such  as  that  described  on  p.  398,  is  converted  to  the 
carbonate  by  means  of  carbon  dioxide  obtained  by  heating  limestone. 

At  present  all  these  methods  are  in  use.  The  great  advantages  of 
the  Solvay  over  the  Leblanc  process  are  to  some  extent  counter- 
balanced by  the  fact  that  the  latter  gives  the  valuable  by-product 
hydrochloric  acid.  It  is  probable  that  in  course  of  time  both  will  be 
superseded  by  the  electrolytic  method. 

Properties  of  Sodium  Carbonate  —  Thesubstanceseparating 
from  solution  at  room  temperature  is  the  decahydrate,  Na2CO3,ioH2O. 
It  has  a  high  tension  of  aqueous  vapour  and  therefore  loses  water 
(effloresces)  on  exposure  to  air,  with  formation  of  the  monohydrate. 
An  intermediate  hydrate,  Na2CO3,7H2O,  can  be  obtained  by  crystal- 
lization from  warm  solutions  under  certain  conditions.  The  deca- 
hydrate and  heptahydrate  are  in  equilibrium  with  the  saturated 
solution  at  32°,  the  heptahydrate  and  monohydrate  at  35.4°.  The 
solubilities  of  the  decahydrate  and  heptahydrate  increase,  that  of  the 
monohydrate  diminishes,  with  rise  of  temperature,  from  which  it 
follows  that  the  salt  must  show  a  maximum  solubility  about  35.4° 
(cf.  sodium  sulphate). 

The  anhydrous  salt  melts  at  852°.  Above  this  temperature  it  begins 
to  decompose,  giving  off  carbon  dioxide. 

The  aqueous  solution  of  sodium  carbonate  is  slightly  alkaline  owing 
to  hydrolysis.  This  has  already  been  fully  explained  (p.  340). 

Sodium  Bicarbonate,  NaHCO3,  is  obtained  by  passing  carbon 
dioxide  into  a  solution  of  sodium  carbonate  : 


As  already  explained,  this  reaction  is  reversible;  even  the  aqueous 
solution  of  sodium  bicarbonate  gives  off  carbon  dioxide  when  boiled. 
The  aqueous  solution  is  practically  neutral  owing  to  the  very  slight 
onisation  of  the  HCO3'  ion.  The  solubility  in  grams  per  100  grams 
of  water  is  as  follows  : 

Temperature      .      o°       10°       20°      30°      40°       50°       60° 
Solubility  .         .      6.9      8.4       9.6      ii.i     12.7      14.5     16.4. 

The  salt  is  used  for  making  baking  powder. 


SODIUM  407 

Sodium  Sulphate — This  salt  occurs  naturally  in  the  salt  de- 
posits in  the  anhydrous  form  as  thenardite  (rhombic  crystals).  It 
is  obtained  on  the  large  scale  in  the  first  stage  of  the  Leblanc  process 
(p.  403)  and  also  in  the  preparation  of  nitric  acid  (p.  223).  The  salt 
which  separates  from  solution  at  room  temperature  is  the  decahydrate, 
Na2SO4,ioH2O,  which  has  been  known  for  centuries  as  Glauber's 
salt.  Above  33°,  the  anhydrous  salt  separates  from  solution. 

The  solubility  curve  of  sodium  sulphate  is  shown  in  Fig.  33,  where 
the  ordinates  represent  solubilities  and  the  abscissae  temperatures. 
The  curve  BC  represents  the  solubility  curve  of  the  decahydrate, 
vhich  increases  with  rise  of  temperature,  and  CD  that  of  the  anhy- 
drous salt,  which  shows  that  the  solubility  diminishes  with  increasing 
temperature.  The  point  C,  at  32.4°,  represents  the  maximum  solubility 
of  the  salt ;  it  is  that  temperature  at  which  decahydrate  and  anhy- 
drous salt  are  in  equilibrium  with  the  solution.  It  must  be  carefully 
remembered  that  the  sharp  change  in  direction  at  C  corresponds 
with  a  change  in  the  solid  from  decahydrate  to  anhydrous  salt  and 
has  nothing  to  do  with  changes  in  the  solution. 

The  readiness  with  which  sodium  sulphate  forms  supersaturated 
solutions  has  already  been  referred  to.  If  such  a  supersaturated 
solution^  cooled  to  5°,  crystals  of  a  heptahydrate,  Na2SO4,7H2O,  are 
obtained.  This  salt  is  unstable  in  contact  with  the  saturated  solution 
under  all  conditions. 

Sodiuri  Nitrate,  NaNO3 — This  salt  occurs  naturally  as  Chili 
saltpetre,  £.nd  is  purified  by  crystallization  from  water.  It  forms 
rhombic  crystals,  which,  unlike  those  of  the  corresponding  potassium 
salt,  are  delquescent.  It  is  largely  used  as  a  manure,1  and  also  for 
the  preparation  of  nitric  acid  and  of  potassium  nitrate  (q-v,}. 

Sodium  JJitrite,  NaNO2,  can  be  prepared  by  strongly  heating 
sodium  nitratt,  or,  better,  by  heating  the  latter  salt  with  metallic  lead, 
and  is  purified  by  crystallization  : 

NaNO3  +  Pb->NaNO2  +  PbO. 

It  is  used  in  the  preparation  of  nitrous  acid. 

Sodium  Phosphates — These  salts  have  already  been  described 
in  connexion  witl  phosphoric  acid  (p.  251).  The  ordinary  phosphate 
of  commerce  is  tie  disodium  salt,  obtained  by  neutralizing  phosphoric 
acid  with  sodium  :arbonate  and  crystallizing.  Below  36°  the  dodeca- 

1  In  1908  over  two  trillion  tons  of  nitrate,  worth  about  ^20,000,000,  were  exported 
from  Chili. 


4o8     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

hydrate,  Na2HPO4,i2H2O,  separates  from  solation  ;  above  this  tem- 
perature the  compound  Na2HPO4,7H0O  is  obtained. 

Sodium  Silicates — Silicates  of  sodium  occur  in  many  rocks. 
A  solution  of  sodium  silicate,  known  technically  as  water-glass,  is 
prepared  by  fusing  two  parts  by  weight  of  sand  or  quartz  with 
one  part  by  weight  of  sodium  carbonate,  dissolving  the  fused  mass 
in  water  and  evaporating  to  a  syrupy  consistency.  The  pure 
silicate,  Na2SiO3,  is  obtained  in  crystalline  form  by  dissolving  the 
calculated  amount  of  freshly  precipitated  silicic  acid  (p.  369)  in 
sodium  hydroxide  solution  and  precipitating  by  means  of  alcohol. 

As  silicic  acid  is  an  extremely  weak  acid,  solutions  of  sodium  silicate 
are  strongly  alkaline  owing  to  hydrolysis. 

Sodium  Sulphides — The  acid  sulphide,  NaHS,  is  obtained  by 
saturating  sodium  hydroxide  solution  with  hydrogen  sulphide  and 
evaporating  to  dryness,  when  the  salt  separates  in  colourless  crystals. 
The  normal  sulphide,  Na2S,  is  obtained  by  adding  to  the  acid  sak  an 
equivalent  of  sodium  hydroxide  and  evaporating  the  solution,  ?vhen 
it  separates  as  the  nonahydrate  in  colourless  crystals : 

NaHS  +  NaOH->Na2S  +  H2O. 

The  behaviour  of  aqueous  solutions  of  these  salts  has  already  been 
explained.  The  acid  sulphide  is  slightly,  the  normal  salt  ?ery  con- 
siderably hydrolyzed  in  aqueous  solution. 

Tests  for  Sodium — As  sodium  salts  are  all  highlyionised  in 
solution,  the  tests  for  the  positive  component  of  the  salty  are  really 
tests  for  sodium  ion.  Since  practically  all  sodium  salts  are  soluble 
in  water,  it  cannot  usefully  be  tested  for  by  precipitation  reactions. 
Sodium  salts  are  always  recognized  by  placing  a  platinum  wire,  to 
which  a  little  of  the  salt  adheres,  in  a  colourless  B/msen  flame, 
when  a  characteristic  intense  yellow  colour  is  obtained  Conclusive 
evidence  as  to  the  presence  of  sodium  is  obtained  by  examining  the 
coloured  flame  with  an  instrument  known  as  the  spectipscope. 

The  Spectroscope.  Spectrum  Analysis — Not  only 
sodium,  but  all  the  alkali  metals,  give  .characteristic  colours  to  the 
Bunsen  flame,  and  can  thus  readily  be  distinguished  with  the  naked 
eye.  When  salts  of  more  than  one  metal  are  present,  however,  this 
method  cannot  be  used,  as  the  more  intense  flame  riasks  the  others. 
By  means  of  the  spectroscope,  however,  the  metals  present  in  a 
mixture  can  be  separately  detected. 

Light  of  a  definite  colour,  for  example  yelloW  is  made  up  of 
vibrations,  whose  wave-lengths  vary  within  nanpw  limits.  White 


SPECTRUM   ANALYSIS  409 

light,  on  the  other  hand,  is  made  up  of  vibrations  of  all  wave-lengths 
within  a  considerable  range.  When  a  beam  of  white  light,  which 
has  passed  through  a  slit,  falls  on  a  triangular  prism  (glass  prisms 
are  generally  used),  the  rays  are  bent  or  refracted  to  a  different 
extent  on  passing  through  the  prism,  those  of  short  wave-length,  e.g. 
violet,  being  most,  and  those  of  longer  wave-length,  e.g.  red,  being 
least  refracted.  Therefore,  if  a  screen  is  placed  behind  the  prism, 
a  so-called  "  continuous  "  spectrum  made  up  of  the  different  colours 
of  the  rainbow  is  seen  ;  it  is  formed  of  a  very  large  number  of  images 
of  the  slit  placed  side  by  side.  The  spectrum  will  also  be  seen  by 
an  observer  properly  placed  behind  the  slit.  The  spectroscope,  an 
instrument  designed  for  the  examination  of  spectra,  consists  essentially 
of  a  prism,  a  tube  provided  with  a  narrow  slit  at  the  end  remote  from 
the  prism,  and  a  movable  telescope,  placed  behind  the  prism,  by 
means  of  which  the  image  of  the  slit  can  be  seen. 

When  a  Bunsen  flame  coloured  by  sodium  is  observed  through 
the  spectroscope,  only  one  yellow  image  of  the  slit  is  seen  against 
a  dark  background.  This  is  due  to  the  fact  that  sodium  light 
is  all  of  one  wave-length,  it  is  therefore  not  scattered  on  passing 
through  the  prism,  and  only  one  image  of  the  slit  is  seen.  Light 
made  up  of  vibrations  of  one  wave-length  only  is  termed  mono- 
chromatic. 

If  a  flame  coloured  by  potassium  is  examined  in  the  same  way,  two 
images  of  the  slit — two  lines,  as  they  are  called — are  to  be  seen,  one 
in  the  region  corresponding  with  violet  in  the  ordinary  spectrum,  the 
other  in  the  red  region.  Lithium  is  characterized  by  two  lines, 
a  well-marked  one  in  red,  and  a  more  feeble  one  in  the  orange.  It 
is  only  in  very  rare  cases  that  lines  belonging  to  two  different 
metals  overlap,  and  therefore  the  constituents  in  a  mixture  of  lithium, 
sodium  and  potassium  can  readily  be  detected  by  means  of  the 
spectroscope. 

One  of  the  great  advantages  of  spectroscopic  analysis  is  that  only 
minute  quantities  of  the  different  substances  are  required  for  the 
purpose  ;  it  is  stated  that  one  ten-millionth  of  a  gram  of  sodium  can 
be  detected  in  this  way.  As  sodium  is  very  widely  distributed  in 
dust,  etc.,  it  is  difficult  to  avoid  the  appearance  of  the  character- 
istic line  in  the  spectrum.  As  each  element  has  its  own  characteristic 
spectrum,  it  can  readily  be  understood  that  many  elements  were  first 
discovered  by  means  of  the  spectroscope.  Thus  the  two  rare  elements 
of  the  alkali  sub-group,  rubidium  and  caesium,  were  detected  by 
Bunsen  and  Kirchhoff  in  the  waters  of  a  mineral  spring  at  Durkheim 


4io    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

(p.  418),  and  helium,  thallium,  scandium,  and  many  other  elements 
were  discovered  in  this  way. 

Line  spectra,  such  as  we  have  been  describing,  are  given  only  by 
substances  in  the  state  of  vapour.  The  spectra  of  metals  are  obtained 
by  volatilizing  the  corresponding  salts  in  a  colourless  flame,  as 
already  indicated.  For  the  same  purpose,  the  permanent  gases  are 
sealed  up  in  glass  tubes,  and  rendered  luminescent  (p.  355)  by  means 
of  an  electric  discharge.  Incandescent  solids,  e.g.  carbon  particles 
in  a  luminous  flame,  give  a  continuous  spectrum. 


POTASSIUM 

Symbol,  K.     Atomic  weight  =39.1.     Molecular  weight  =39.1  (probably). 

Occurrence  —  Potassium  occurs  in  nature  mainly  in  the  form  of 
silicates  in  rocks,  especially  in  felspar  and  mica.  When  these  rocks 
are  broken  up,  the  soluble  potassium  salts  pass  into  the  soil,  from 
which  they  are  taken  up  by  plants.  Potassium  salts  are  essential  for 
the  growth  of  plants,  and  are  taken  up  in  much  larger  proportion 
than  sodium  salts.  When  plants  are  burned  the  potassium  is  found 
in  the  ash  as  carbonate,  and  this  formerly  represented  the  chief 
commercial  source  of  potassium  compounds. 

•Part  of  the  potassium  salts  from  the  disintegration  of  rocks  finds 
its  way  into  seas,  lakes  and  mineral  springs.  When  an  inland  sea 
evaporates,  the  less  soluble  sodium  chloride  separates  out  first,  and 
above  this,  at  a  later  stage,  are  deposited  the  more  soluble  salts, 
chiefly  those  of  potassium  and  magnesium.  The  Stassfurt  deposits, 
formed  in  this  way,  now  constitute  the  chief  source  of  potassium, 
the  more  important  compounds  being  carnallite,  MgCl2,KCl,6H2O, 
kainite,  MgSO4,KCl,3H2O,  and  sylvine,  KC1. 

Preparation  of  Metal  —  Like  sodium,  potassium  was  first 
obtained  by  Davy  (1807)  by  electrolysis  of  moist  potassium  hy- 
droxide. 

The  small  quantities  of  this  metal  required  commercially  are  pre- 
pared by  chemical  methods.  The  chief  method  consists  in  strongly 
heating  a  mixture  of  potassium  carbonate  and  finely  divided  carbon  : 


Potassium  and  carbon  monoxide  combine  under  certain  conditions 
to  form  potassium  carboxide,  K6(CO)6,  a  highly  explosive  compound, 
and  in  order  to  avoid  this  the  metal  must  be  rapidly  condensed, 


POTASSIUM  411 

best  by  using  flat  receivers  surrounded  by  cold  water.  The  metal 
is  collected  under  mineral  oil. 

The  intimate  mixture  of  carbonate  and  carbon  is  usually  obtained 
by  raising  a  mixture  of  the  carbonate  and  tar  to  a  low  red  heat. 

Electrolytic  methods  could  also  be  used  for  preparing  potassium  if 
the  metal  were  of  sufficient  commercial  importance. 

Properties — In  all  its  properties  potassium  closely  resembles 
sodium.  It  is  a  soft,  silvery-white  metal,  which  can  be  obtained  in 
cubic  crystals  by  sublimation  ;  it  melts  at  62°,  and  boils  at  757°. 
The  vapour  a  little  above  the  melting-point  is  greenish,  at  higher 
temperatures  it  becomes  violet.  It  is  probably  monatomic  in  the 
state  of  vapour. 

Potassium  oxidizes  in  moist  air  or  oxygen,  and  combines  with  the 
halogens  even  more  vigorously  than  sodium  does.  It  reacts  vigorously 
with  water : 

2K  +  2H2O->2KOH  +  H2, 

and  so  much  heat  is  given  out  in  the  process  that  the  evolved 
hydrogen  catches  fire  and  burns  with  a  lavender  flame. 

Potassium  Hydride,  KH,  resembles  sodium  hydride  in  its 
mode  of  preparation  and  properties. 

Potassium  Oxides — At  least  two  oxides  of  potassium,  the 
monoxide,  K2O,  and  a  peroxide,  K2O4,  are  definitely  known. 

Potassium  Monoxide,  K2O,  is  obtained  by  incomplete  com- 
bustion of  the  metal  in  dry  oxygen,  the  excess  of  metal  being  removed 
by  distillation  in  a  vacuum.  It  occurs  in  yellowish-wliite  crystals,  and 
combines  very  vigorously  with  water. 

Potassium  Peroxide,  K2O4,  is  the  chief  product  obtained 
when  potassium  is  burned  with  free  access  of  air.  It  is  a  yellow 
powder,  which  is  decomposed  by  water,  with  formation  of  potassium 
hydroxide,  hydrogen  peroxide,  and  oxygen  : 


Potassium  Hydroxide,  KOH— Preparation— (i)  By  boiling 
potassium  carbonate  with  calcium  hydroxide  or  by  the  action  of  the 
latter  compound  on  potassium  sulphate  : 

+  2KOH 
+  2KOH. 

The  details  have  already  been  given  under  sodium  hydroxide. 

(2)  As  in  the  case  of  sodium  hydroxide,  by  electrolysis  of  an  aqueous 
solution  of  potassium  chloride. 


412     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

Properties  —  Potassium  hydroxide  is  a  white  amorphous  substance, 
which  melts  at  360°.  It  is  usually  sold  in  sticks.  Three  hydrates, 
KOH,H20  ;  KOH,2H2O  and  KOH,4H2O,  are  definitely  known.  In 
all  its  properties  it  closely  resembles  sodium  hydroxide.  When  solu- 
tions of  potassium  or  sodium  hydroxide  are  evaporated  by  boiling 
the  temperature  gradually  rises  without  the  separation  of  solid  at 
any  stage,  and  finally  the  fused  compound  is  obtained.  This 
curious  behaviour  depends  on  the  fact  that  at  no  stage  during  the 
evaporation  is  there  a  saturated  solution,  whose  vapour  pressure 
reaches  one  atmosphere. 

Potassium  Chloride,  KC1  —  This  salt  occurs  in  the  Stassfurt 
deposits  as  sylvine  and  as  a  constituent  of  carnallite,  MgCl2,KCl, 
6H2O.  It  can  be  prepared  by  the  general  methods  for  a  soluble 
salt  (p.  390),  and  commercially  is  chiefly  obtained  from  carnallite. 
When  the  latter  compound  is  dissolved  in  water,  it  splits  up  into  its 
components,  and  on  evaporating  the  solution  the  potassium  salt,  being 
least  soluble,  separates  out  first. 

Properties  —  The  salt  crystallizes  in  the  anhydrous  form  in  cubic 
crystals.  No  hydrates  of  potassium  chloride,  bromide  or  iodide  are 
known.  Unlike  sodium  chloride,  the  solubility  of  potassium  chloride 
in  water  increases  considerably  with  rise  of  temperature. 

Potassium  Bromide,  KBr  —  The  salt  can  be  prepared  by  the 
general  methods,  but  is  almost  invariably  obtained  by  adding  bromine 
to  a  hot  aqueous  solution  of  potassium  hydroxide  as  long  as  the  colour 
disappears  ;  a  mixture  of  bromide  and  bromate  is  thus  obtained  • 


If  both  bromide  and  bromate  are  required,  the  salts  are  separated  by 
fractional  crystallization,  the  bromate  being  least  soluble.  If  only  the 
bromide  is  wanted  the  solution  is  evaporated  to  dryness,  the  residue 
mixed  with  charcoal,  strongly  heated  to  reduce  the  bromate  to  bromide  : 


and  the  potassium  bromide  purified  by  recrystallization. 

Properties  —  The  salt  separates  from  solution  in  colourless,  anhy- 
drous cubic  crystals.  It  is  used  in  medicine,  and  also  in  photography 
for  preparing  silver  bromide  (^.;z>.). 

Potassium  Iodide,  KI,  is  prepared  by  the  method  described 
under  potassium  bromide,  iodine  being  substituted  for  bromine,  and 
also  as  follows.  Iron  and  iodine  are  brought  together  below  water, 
when  they  combine,  forming  a  solution  of  the  compound  Fe3I8  ;  the 


POTASSIUM  413 

latter  is  then  decomposed  by  potassium  hydroxide  with  formation 
of  the  compound  Fe3O4,  which  is  precipitated,  and  potassium  iodide, 
which  remains  in  solution  : 

Fe3I8+8KOH->Fe304  +  8KI+4H20. 

Potassium  iodide  occurs  in  cubic  crystals,  which  are  very  soluble  in 
water.  It  is  used  in  medicine  and  in  photography. 

Potassium  Chlorate,  Y£\QZ—  Preparation—  (\)  By  the  action 
of  chlorine  in  excess  on  a  hot  solution  of  potassium  hydroxide  l 
(p.  181): 


As  potassium  chlorate  is  not  very  soluble  in  cold  water,  the  salts  can 
readily  be  separated  by  crystallization. 

(2)  By  the  above  method  only  one-sixth  of  the  potassium  appears 
as  chlorate,  and  it  is  much  more  economical  first  to  prepare  calcium 
chlorate  by  the  action  of  chlorine  on  hot  milk  of  lime  : 


The  proper  amount  of  potassium  chloride  is  then  added  and  the 
solution  evaporated  to  a  definite  density  and  set  aside,  when  potas- 
sium chlorate  separates  out. 

(3)  Potassium  chlorate  is  now  largely  prepared  by  electrolytic 
methods,  for  example,  by  electrolysis  of  a  hot  solution  of  potassium 
chloride,  the  chlorine  and  alkali  formed  at  the  anode  and  cathode 
respectively  reacting  to  form  the  chlorate,  which  separates  out  on 
concentrating  the  solution. 

Properties  —  Potassium  chlorate  occurs  in  anhydrous  monoclinic 
leaflets,  and  melts  at  370°.  The  solubility  in  water  (grams  in  100 
grams  of  water)  is  given  in  the  accompanying  table  : 

Temperature      o°       10°      20°      30°        40°       50°      70°       100° 
Solubility  3.3       5.0      7.1      10.1        14.5      19.7      32.5      56.0 

When  heated  a  little  above  its  melting-point,  it  rapidly  gives  off 
oxygen,  and  finally  only  the  chloride  remains.  If  the  reaction  is 
stopped  at  an  intermediate  point,  the  residue  is  found  to  contain 
potassium  perchlorate,  KC1O4,  which  can  readily  be  separated 

1  It  has  recently  been  shown  that  excess  of  chlorine  (which  produces  free 
hypochlorous  acid)  is  much  more  important  than  high  temperature  for  the 
production  of  alkali  chlorate  from  hypochlorite. 


414     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

from  the  chloride  and  unaltered  chlorate  by  means  of  its  sligh 
solubility  in  water. 

Potassium  chlorate  is  used  as  a  source  of  oxygen,  as  well  as  in  the 
firework  and  match  industry  (p.  242),  and  is  also  made  use  of  in 
medicine. 

Potassium  Perchlorate,  KC1O4,  prepared  from  the  chlorate 
as  just  described,  forms  colourless  rhombic  crystals.  At  o°,  25°  and 
50°  100  grains  of  water  dissolve  0.71,  1.96  and  5.34  grams  of  the  salt 
respectively.  It  is  a  less  powerful  oxidizing  agent  than  the  chlorate  ; 
it  begins  to  give  off  oxygen  when  the  temperature  exceeds  400°. 

Potassium  Carbonate,  K2CO3—  This  salt  was  formerly  pre- 
pared by  extracting  the  ash  of  plants  with  water  and  evaporating,  but 
since  the  discovery  of  the  Stassfurt  deposits  is  prepared  from  the 
chloride  by  a  method  analogous  to  the  Leblanc  process  ;  or  by 
electrolysis,  and  treatment  of  the  hydroxide  with  carbon  dioxide. 
On  account  of  the  considerable  solubility  of  potassium  bicarbonate  in 
water,  a  method  of  preparation  'analogous  to  the  ammonia-soda 
process  is  not  workable. 

Potassium  carbonate  is  a  colourless  salt,  extremely  soluble  in  water. 
When  the  anhydrous  salt  is  exposed  to  the  air  it  takes  up  moisture 
and  becomes  liquid  (deliquesces),  owing  to  the  fact  that  the  vapour 
pressure  of  its  saturated  solution  is  less  than  the  average  pressure  of 
aqueous  vapour  in  the  atmosphere  (p.  403).  It  appears  to  form  a 
number  of  hydrates  with  water,  among  others  a  dihydrate,  but  the 
lacts  have  not  yet  been  clearly  established.  The  aqueous  solution 
of  the  salt  is  alkaline  owing  to  hydrolysis  (p.  267)  : 

K2CO3+HOH^KHCO3+KOH. 

Owing  to  its  great  affinity  for  water,  the  anhydrous  salt  is  used  as  a 
drying  agent. 

Potassium  Bicarbonate,  KHCO3—  This  salt  can  be  pre- 
pared by  passing  carbon  dioxide  into  a  concentrated  solution  of  the 
normal  carbonate  ;  being  less  soluble  than  the  latter,  it  separates 
from  solution  and  can  be  purified  by  recrystallization  : 


Potassium  bicarbonate  is  more  soluble  in  water  than  the  correspond- 
ing sodium  salt. 

The  aqueous  solution  of  the  pure  salt  is  practically  neutral  owing 
to  the  very  slight  ionisation  of  the  HCO3'  ion  (p.  340). 

Potassium  Sulphate,  K2SO4—  This  salt  occurs  in  the  Stass- 


POTASSIUM*  415 

furt  deposits  chiefly  as  kainite,  MgSO4,KCl,3H.>O,  and  polyhalite, 
MgSO4,2CaSO4,K2SO4,2H2O.  It  is  obtained  from  kainite  by  treating 
with  cold  water,  when  a  separation  into  the  difficultly  soluble  schonite, 
MgSO4,K2SO4,6H2O,  and  the  readily  soluble  magnesium  chloride 
occurs.  The  latter  is  completely  removed,  and  the  schonite  treated 
with  excess  of  potassium  chloride,  when  potassium  sulphate  crystal- 
lizes out  : 


Potassium  sulphate  is  also  prepared  by  heating  the  chloride  with 
concentrated  sulphuric  acid. 

Potassium  sulphate  occurs  in  rhombic  prisms,  and,  in  contrast  to 
sodium  sulphate,  forms  no  hydrates.  About  i  part  of  the  salt  dis- 
solves in  10  parts  of  water  at  15°.  It  finds  application  as  a  manure. 

Potassium  Hydrogen  Sulphate  (potassium  bisul- 
phate),  KHSO4  —  This  salt  is  prepared  by  heating  potassium 
chloride  with  excess  of  sulphuric  acid  : 

KC1  +  H2SO4->KHSO4  +  HC1. 

It  forms  monoclinic  prisms  which  melt  about  210°,  and  on  further 
heating  break  up  into  the  pyrosulphate  and  water  (p.  312). 

2KHSO4->K2S2O7+  H2O. 

It  is  extremely  soluble  in  water,  and  the  solution  is  strongly  acid 
owing  to  the  fact  that  the  anion,  HSO4',  splits  off  H'  ions  (p.  310). 

Potassium  Nitrate  —  This  salt  has  been  known  from  the  earliest 
times.  It  is  formed  when  nitrogenous  organic  matter,  ammoniacal 
compounds,  dung,  etc.,  decay  in  contact  with  potassium  salts,  con- 
tained, for  example,  in  wood  ashes,  and  therefore  the  conditions  were 
favourable  for  its  formation  in  the  neighbourhood  of  dwellings  at  a 
very  early  stage  of  civilization.  The  oxidation  in  this  case  is  accom- 
plished by  the  oxygen  of  the  air  with  the  help  of  certain  bacteria,  and 
does  not  take  place  in  the  absence  of  the  bacteria.  This  process  is 
used  commercially  in  the  so-called  "  nitre  plantations."  Dung  and 
other  animal  refuse  are  mixed  with  wood-ashes  and  exposed  in  heaps 
to  the  action  of  the  air.  After  a  year  or  two  the  potassium  nitrate  is 
washed  out  and  purified  by  recrystallization. 

Potassium  nitrate  is  now  usually  prepared  by  mixing  hot  saturated 
solutions  of  sodium  nitrate  and  potassium  chloride.  In  such  a  mixture 
the  four  ions  K',  Na',  Cl'  and  NO3'  are  present,  and  the  salt,  whose 
solubility  product  (p.  439)  is  first  reached  under  definite  conditions, 
will  deposit  from  solution.  A  glance  at  the  solubility  table  shows 


4i6    A   TEXT-BOOK  "OF   INORGANIC   CHEMISTRY 

that  at  high  temperatures  sodium  chloride  is  the  least  soluble  salt 
which  can  be  formed  from  the  ions  ;  it  therefore  deposits  in  consider- 
able amount.  When  the  solution  is  then  allowed  to  cool,  potassium 
nitrate  becomes  the  least  soluble  salt,  and  it  separates  in  a  fairly  pure 
condition.  It  is  then  purified  by  recrystallization.  It  is  evident  that 
the  purification  of  a  salt  by  crystallization  is  most  readily  accom- 
plished when,  as  in  the  present  case,  the  solubility  changes  greatly 
with  change  of  temperature. 

Properties  —  Potassium  nitrate  is  a  dimorphous  substance  ;  at 
low  temperatures  it  separates  from  solution  in  rhombic,  at  high 
temperatures  in  rhombohedral  crystals.  It  melts  at  339°,  and  at 
higher  temperatures  loses  oxygen  and  forms  potassium  nitrite  : 

2KNO3->2KNO2  +  O2. 

The  fused  nitrate  has  powerful  oxidizing  properties.  Sulphur  (or 
charcoal)  thrown  into  a  crucible  containing  it  catches  fire  and  burns 
vigorously.  The  salt  is  largely  used  in  the  preparation  of  gunpowder  ; 
it  has  the  advantage  over  sodium  nitrate  of  not  being  deliquescent. 

Gunpowder  —  Gunpowder  is  a  mixture  of  potassium  nitrate, 
sulphur  and  charcoal.  The  proportions  per  cent,  of  the  ingredients 
are  not  quite  constant,  but  are  approximately  as  follows  :  Potassium 
Nitrate,  75;  Sulphur,  10;  Charcoal,  15.  On  this  basis  we  might 
expect  the  decomposition  to  be  represented  by  the  equation 


As  a  matter  of  fact,  the  reactions  are  much  more  complex.  Besides 
potassium  sulphide,  the  solid  products  include  potassium  sulphate, 
potassium  thiosulphate  and  other  salts  ;  and  among  the  gases,  besides 
carbon  dioxide  and  nitrogen,  carbon  monoxide  and  hydrogen  sulphide 
are  formed.  About  42  to  45  per  cent,  goes  off  as  gas,  and  55  to  58 
per  cent,  of  solid  matter  remains. 

The  efficiency  of  the  powder  depends  upon  the  almost  instantaneous 
production  in  a  small  space  of  a  considerable  quantity  of  gas,  which 
therefore  exerts  a  very  high  pressure,  and  this  pressure  is  still  further 
increased  by  the  heat  given  out  in  the  combustion  of  the  powder.  It 
has  been  found  that  the  pressure  of  the  gaseous  products  reaches 
6400  atmospheres,  and  the  temperature  reaches  2000°. 

Potassium  Sulphides  —  The  normal  and  acid  sulphides,  K2S 
and  KHS,  are  prepared  as  described  in  connexion  with  the  corre- 
sponding sodium  salts.  The  former  separates  from  solution  as  the 
pentahydrate,  K2S,5H2O,  in  reddish  crystals. 


LITHIUM  417 

Poly  sulphides )  e.g.  K2S2,  K2S3,  K2S4  and  K2S5  are  also  known. 
The  pentasulphide  is  obtained,  along  with  potassium  thiosulphate,  by 
heating  potassium  carbonate  with  excess  of  sulphur  to  250°  for  some 
time  in  a  closed  vessel : 

3K2CO3  +  6S2->2K2S5+K2S2O3+3CO2. 

The  mixture  of  these  two  compounds  was  formerly  called  liver  of 
sulphur. 

Tests  for  Potassium — As  already  mentioned,  compounds  of 
potassium  are  recognized  by  the  characteristic  lavender  colour  im- 
parted to  a  colourless  flame.  Like  the  sodium  salts,  the  potassium 
salts  are,  almost  without  exception,  readily  soluble  in  water.  The 
exceptions  are  potassium  platinic  chloride,  PtCl4,2KCl,  and  potassium 
acid  tartrate,  which  are  soluble  with  difficulty  in  water.  Potassium 
compounds  alone,  or  in  the  presence  of  sodium  salts,  can  therefore  be 
detected  by  adding  platinic  chloride,  concentrated  hydrochloric  acid 
and  excess  of  alcohol,  when  the  double  salt  separates  in  yellow 
crystals,  or  by  adding  to  the  solution  excess  of  tartaric  acid,  and 
shaking,  when  the  acid  tartrate  separates  out  as  a  crystalline  pre- 
cipitate. It  may  be  mentioned  that  ammonium  salts  give  similar 
precipitates  under  the  same  conditions. 

LITHIUM 
Symbol,  Li.     Atomic  weight=7.o3. 

Occurrence— In  contrast  to  the  compounds  of  sodium  and  potassium,  lithium 
compounds  are  found  very  sparingly  in  nature.  It  occurs,  usually  in  very  small 
proportion,  in  a  number  of  silicates.  Those  richest  in  the  metal  are:  spodumene 
and  petalite,  both  lithium-aluminium  silicates,  the  former  containing  about  3.8 
per  cent,  and  the  latter  about  2  per  cent,  of  lithium,  and  lepidolite  or  lithium 
mica,  an  alkali  aluminium  fluorosilicate  which  contains  0.8  to  2.7  per  cent,  of 
lithium.  It  also  occurs  in  small  quantity  in  certain  mineral  springs,  and  in  the 
ash  of  certain  plants,  such  as  tobacco. 

Preparation  of  Metal— The  metal  is  usually  obtained  by  electrolysis  of  the 
fused  chloride  ;,the  salt  is  heated  to  fusion  in  a  porcelain  crucible,  and  subjected 
to  electrolysis,  a  carbon  rod  being  used  as  anode  and  an  iron  wire  as  cathode. 
Instead  of  the  chloride,  a  mixture  of  lithium  bromide  with  10  to  15  per  cent,  of  the 
chloride,  which  fuses  at  a  lower  temperature  than  the  chloride,  may  be  used. 
The  metal  may  also  be  prepared  by  electrolysis  of  a  solution  of  the  chloride  in 
pyridine. 

Properties— Lithium  is  a  silvery-white  metal,  harder  than  sodium,  but  softer 
than  lead.  It  is  the  lightest  solid  known,  its  density  being  only  0.534  ;  it  floats  on 
mineral  naphtha  ;  it  melts  at  186°. 

Lithium  rapidly  tarnishes  in  moist  air  ;  when  heated  above  200°  it  burns  in  air 
with  a  white  flame  like  that  of  magnesium,  forming  lithium  oxide,  Li2O.  It  acts 
27 


4i8    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

on  water  at  the  ordinary  temperature,  but  much  less  vigorously  than  the  other 
alkali  metals,  lithium  hydroxide,  LiOH,  being  formed  and  hydrogen  set  free. 

Lithium  Oxide,  Li2O,  is  prepared  by  burning  lithium  in  air  or  by  heating  lithium 
nitrate.  It  dissolves  in  water,  forming  lithium  hydroxide,  LiOH.  The  latter 
compound  can  also  be  prepared  by  boiling  lithium  carbonate  with  calcium 
hydroxide  (cf.  sodium  hydroxide,  p.  398).  Lithium  hydroxide  is  a  strong  base. 

Lithium  Chloride,  LiCl,  is  very  soluble  in  water.  Between  o°  and  20°  the 
dihydrate,  LiCl,2H2O,  from  20°  to  100°  the  monohydrate,  and  above  100°  the 
anhydrous  salt  separate  from  solution. 

Lithium  Fluoride,  LiF,  differs  from  the  fluorides  of  the  other  alkali  metals  in 
being  very  slightly  soluble  in  water  (about  i  in  400  at  18°). 

Lithium  Carbonate,  Li2CO3,  is  also  much  less  soluble  in  water  than  the 
other  alkali  carbonates  ;  100  grams  of  water  dissolve  at  o°  1.54  parts,  and  at  20° 
1.33  parts  of  the  salt. 

Lithium  Phosphate,  Li3PO4,  separates  from  solution  as  a  white  crystalline 
precipitate  when  ammonia  and  sodium  phosphate  are  added  to  a  solution  of  a 
lithium  salt.  It  is  very  slightly  soluble  in  water  (i  part  in  2540  parts  of  water  at 
18°),  and  this  is  taken  advantage  of  as  a  test  for  lithium  compounds  and  in  sepa- 
rating it  from  mixtures. 

Tests  for  Lithium— Apart  from  the  behaviour  of  the  phosphate,  lithium  is 
characterized  by  the  crimson  colour  it  imparts  to  flame,  and  by  the  presence  of  a 
red  and  a  yellow  line  in  the  spectrum  (p.  409). 

RUBIDIUM  AND  CAESIUM 

These  rare  elements  were  discovered  by  Bunsen  and  Kirchhoff  (1860-1861)  in 
the  mineral  waters  of  Durkheim  by  means  of  the  spectroscope.  They  also  occur 
in  certain  lepidolites  (p.  417),  and  rubidium  is  now  obtained  almost  entirely  from 
the  Stassfurt  deposits,  in  which  it  occurs  in  very  small  amount  as  rubidium  car- 
nallite,  MgCl2,RbCl.  It  is  separated  from  the  other  alkali  metals  by  taking 
advantage  of  the  comparative  insolubility  of  rubidium  alum,  AySO^.RbgSO^ 
24H2O.  The  only  naturally  occurring  substance  rich  in  caesium  is  the  rare 
mineral pollux,  a  silicate  of  aluminium  and  caesium,  which  contains  up  to  32  per 
cent,  of  the  latter  element. 

Both  are  soft,  silvery-white  metals,  which  ignite  at  the  ordinary  temperature  in 
moist  air  and  also  in  dry  oxygen.  Rubidium  melts  at  38°,  caesium  at  26.5°. 

The  chemistry  of  rubidium  and  caesium  is  practically  identical  with  that  of 
potassium.  The  only  outstanding  difference  is  that  the  former  elements  give 
polyhalogen  compounds  containing  three  and  five  atoms  respectively  of  the  halo- 
gens, e.g.  RbBr3,CsBr3,CsBr6,  in  addition  to  the  normal  compounds  with  one  atom 
of  halogen. 

AMMONIUM  SALTS 

It  has  already  been  mentioned  that  the  salts  containing  the  NH4or 
ammonium  group  show  a  remarkable  analogy  with  those  of  the 
alkali  metals,  especially  potassium,  and  it  is  therefore  convenient  to 
deal  with  them  here. 

The  chief  source  of  ammonia  and  its  compounds  is  the  ammoniacal 


AMMONIUM  COMPOUNDS  419 

liquor  of  the  gas-works  (p.  214).  When  steam  is  blown  through  the 
liquor  the  easily  hydrolyzable  salts  (carbonate  and  sulphide)  are  de- 
composed and  the  ammonia  passes  off.  The  stable  salts  remaining 
in  the  liquor  are  decomposed  by  boiling  with  milk  of  lime.  The 
process  can  be  made  continuous  by  using  a  rectifying  column  on  the 
top  of  a  still,  the  easily  hydrolyzable  salts  being  decomposed  by  steam 
in  the  column,  the  stable  salts  by  lime  in  the  still.  The  ammonia  is 
absorbed  in  hydrochloric  acid.  The  solution  is  evaporated,  and  the 
solid  ammonium  chloride  thus  obtained  purified  by  sublimation.  From 
the  chloride  all  the  other  ammonium  compounds  can  readily  be  obtained 
Ammonium  Chloride  (sal  ammoniac),  NH4C1,  prepared 
as  mentioned  above,  usually  occurs  in  feathery  groups  of  small  octa- 
hedral (more  rarely  cubic)  crystals,  which  have  a  salt  taste.  When 
heated  it  sublimes  before  the  melting-point  is  reached,  and  the  vapour 
appears  (from  vapour  density  measurements)  to  be  completely  dis- 
sociated into  ammonia  and  hydrogen  chloride  : 

NH4C1^NH3+HC1. 

Baker  has  shown  that  neither  of  the  above  reversible  reactions  takes 
place  when  the  substances  are  perfectly  dry  ;  ammonium  chloride 
does  not  dissociate,  nor  is  the  chloride  formed  when  ammonia  and 
hydrogen  chloride  are  mixed. 

At  o°,  20°,  and  40°  water  dissolves  29.7,  37.2,  and  45.8  grams  respec- 
tively of  the  chloride.  As  ammonia  is  a  weak  base  (see  below)  the 
chloride  is  slightly  hydrolyzed  in  aqueous  solution  : 


and  this  is  the  more  pronounced  the  higher  the  temperature.  Further, 
as  the  hydroxide  in  aqueous  solution  is  partly  dissociated  into  water 
and  ammonia,  the  latter  escapes  when  the  solution  is  boiled,  so  that 
the  aqueous  solution  of  ammonium  chloride  (and  of  all  other  ammo- 
nium salts),  which  is  very  faintly  acid  under  ordinary  conditions, 
becomes  much  more  acid  on  boiling. 

Ammonium  Hydroxide,  NH4OH  —  When  ammonia  gas  is 
passed  into  water  it  is  readily  absorbed,  and  the  resulting  strongly 
alkaline  solution  undoubtedly  contains  ammonium  hydroxide,  NH4OH, 
which  is  largely  ionised  into  NH4*  and  OH'  ions.  There  is  evidence, 
however,  that  the  combination  with  water  is  not  complete  ;  in  other 
words,  the  solution  contains  free  ammonia.  The  state  of  affairs  may 
therefore  be  represented  by  the  equations 


420    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

At  present  no  quite  satisfactory  method  is  known  by  means  of  which 
the  relative  amounts  of  these  substances  can  be  determined. 

Ammonium  Sulphate,  (NH4)2SO4  —  The  crude  salt  is  obtained 
by  absorbing  the  ammonia  from  gas  liquors  in  sulphuric  acid,  and  is 
purified  by  recrystallization.  It  separates  from  solution  in  colourless, 
anhydrous,  rhombic  crystals,  isomorphous  with  potassium  sulphate. 
The  crude  salt  is  largely  used  as  a  manure. 

Ammonium  Carbonates  —  Commercial  ammonium  carbonate, 
also  known  as  ammonium  sesquicarbonate^  is  obtained  by  strongly 
heating  in  a  retort  a  mixture  of  ammonium  chloride  and  calcium 
carbonate  and  condensing  the  vapours  in  a  receiver.  It  is  purified 
by  sublimation.  The  substance  is  a  loose  combination  of  a  molecule 
of  ammonium  acid  carbonate  and  a  molecule  of  ammonium  carbamate 
NH4NH2CO2  (that  is,  a  molecule  of  the  normal  carbonate  minus 
I  H2O),  and  therefore  has  the  formula,  NH4HCO3,NH4NH2CO2. 
The  compound  usually  occurs  in  crystalline  masses.  It  loses  am- 
monia at  the  ordinary  temperature,  for  which  reason  it  smells  strongly 
of  the  gas,  and  also  carbon  dioxide,  and  finally  only  the  bicarbonate 
remains.  On  heating  above  60°  it  decomposes  completely  into  water, 
ammonia  and  carbon  dioxide. 

When  the  sesquicarbonate  is  dissolved  in  water  the  carbamate 
combines  with  the  latter  to  form  the  normal  carbonate,  (NH4)2CO3,  so 
that  the  solution  contains  a  mixture  of  normal  and  acid  carbonate. 
Finally,  when  sufficient  ammonia  is  added  to  convert  the  acid  car- 
bonate into  the  normal  carbonate,  a  solution  containing  the  latter  salt 
only  is  obtained  : 


(1)  NH4HC03,NH4NH2C02+H2O^NH4HC03 

(2)  NH4HC03+(NH4)2C03+  NH3->2(NH4)2CO3. 

The  normal  carbonate,  (NH4)2CO3,  separates  in  the  solid  form  when 
ammonia  gas  is  passed  into  a  concentrated  solution  of  the  commercial 
carbonate.  It  gives  up  ammonia  at  room  temperature,  forming  the 
bicarbonate,  and  when  heated  to  60°  decomposes  completely  into 
carbon  dioxide,  ammonia  and  water. 

The  acid  carbonate^  NH4HCO3,  is  obtained  in  the  solid  form  by 
decomposition  of  the  normal  carbonate  or  commercial  carbonate,  as 
stated  above,  and  also  by  passing  carbon  dioxide  into  an  aqueous 
solution  of  the  commercial  carbonate  and  evaporating  : 

NH4HC03,(NH4)2C03+C02+H20->3NH4HCO,» 
The  dry  salt  gives  up  practically  no  ammonia  at  room  temperature. 


AMMONIUM   COMPOUNDS  421 

At  12.5°  100  grams  of  water  dissolve  17.1  grams,  at  21°  21.6  grams  of 
the  salt. 

In  aqueous  solution  the  salts  are  all  considerably  hydrolyzed,  as 
both  base  and  acid  are  weak. 

Ammonium  Sulphides  —  When  hydrogen  sulphide  is  passed 
in  excess  into  a  solution  of  ammonia,  ammonium  hydrosulphide, 
NH4HS,  is  obtained  ;  and  if  an  equivalent  amount  of  ammonia  is 
added  it  might  be  assumed  that  the  normal  sulphide,  (NH4)2S,  is 
formed.  As,  however,  both  base  and  acid  are  weak,  the  latter  com- 
pound is  almost  completely  hydrolyzed  in  solution  : 


Even  the  hydrosulphide  is  hydrolyzed  to  some  extent  under  ordinary 
conditions. 

The  hydrosulphide  can  be  obtained  in  rhombic  leaflets,  stable  at 
room  temperature  but  dissociating  into  ammonia  and  hydrogen 
sulphide  on  heating.  The  normal  sulphide  cannot  be  obtained  at 
room  temperature  in  the  solid  form. 

Solutions  which  contain  only  the  acid  or  normal  sulphide  are 
practically  colourless.  On  standing,  the  hydrogen  sulphide  formed 
by  hydrolysis  in  the  solution  of  the  normal  sulphide  is  slowly  oxidized 
by  atmospheric  oxygen  to  water  and  sulphur,  and  the  latter  then 
combines  with  the  normal  sulphide  to  form  yellow  poly  sulphides,  e.g. 
(NH4)2S4,  (NH4)2S6  and  (NH4)2S7.  The  importance  in  analysis  of 
these  solutions  of  "  yellow  "  ammonium  sulphide  is  referred  to  later 
(pp.  519). 

Ammonium,  NH4  —  As  the  group  NH4  behaves  in  many  respects 
like  an  alkali  metal  it  is  interesting  to  inquire  whether  it  has  been 
isolated.  In  spite  of  many  attempts,  this  has  so  far  not  been  done. 
When  sodium  amalgam  is  added  to  a  strong  solution  of  ammonium 
\  chloride  a  porous,  bulky  mass  is  obtained,  which  breaks  up  into 
mercury,  ammonia,  and  hydrogen.  It  was  formerly  supposed  that 
this  bulky  mass  was  "ammonium  amalgam,"  a  solution  of  NH4  in 
mercury,  but  no  conclusive  evidence  on  this  point  has  been  obtained. 
The  same  amalgam  is  obtained  when  an  aqueous  solution  of  an 
ammonium  salt  is  electrolyzed  with  a  mercury  cathode.  A  further 
attempt  to  obtain  ammonium  by  the  electrolysis  at  -  95°  of  ammonium 
iodide  dissolved  in  liquid  ammonia  was  also  unsuccessful. 

Tests  for  Ammonium  —  Ammonium  compounds  are  easily 
distinguished  by  the  characteristic  odour  of  ammonia  when  they  are 
warmed  with  sodium  hydroxide. 


422     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


COMPARISON  OF  THE  ALKALI  METALS  AND  SUMMARY  > 

The  main  characteristics  of  the  members  of  this  group  are  that  all 
are  soft  metals,  which  readily  oxidize  in  the  air,  and  decompose  water 
at  room  temperature.  They  are  the  most  electro-positive  of  all  the 
metals  ;  the  hydroxides  are  very  strong  bases  ;  practically  all  the  salts 
are  readily  soluble  in  water.  They  impart  characteristic  colours  to 
the  Bunsen  flame.  They  are  all  univalent  elements. 

As  in  other  groups,  the  physical  and  chemical  properties  vary 
regularly  with  increasing  atomic  weight.  This  is  shown,  for  the 
more  important  physical  properties,  in  the  accompanying  table  : — 


Li. 

Na. 

K. 

Rb. 

Cs. 

Atomic  weight 
Density  (solid) 

6.94 
0-534 

23.00 
0.9741 

39.10 
0.863 

8545 
1.52 

132.81 

1.87 

Melting-point  . 

1  86° 

97° 

62.5° 

38-5° 

26.5° 

Boiling-point   . 

red  heat 

878° 

758° 

696° 

670° 

Atomic  volume 

I3-I 

23-7 

45-4 

55-8 

71 

As  regards  chemical  behaviour,  we  have  seen  that  the  activity 
increases  with  increasing  atomic  weight.  This  is  shown  in  combining 
with  oxygen,  rubidium  and  caesium  bursting  into  flame  in  dry 
oxygen  at  room  temperature,  in  the  behaviour  towards  water,  and  in 
the  heat  of  formation  of  their  compounds  ;  also  by  the  fact  that  the 
stability  of  the  bicarbonates  increases  with  increasing  atomic  weight. 

There  are  also  certain  characteristic  differences.  Lithium  differs 
in  several  respects  from  the  other  four,  as  is  not  uncommon  in  the 
case  of  the  first  member  of  a  family.  Thus  lithium  carbonate  and 
phosphate  (and  also  the  fluoride)  are  much  less  soluble  in  water  than 
the  corresponding  salts  of  the  other  alkali  metals  :  properties  which 
recall  the  alkaline  earth  metals.  The  salts  of  sodium  are  more 
generally  soluble  in  water  than  those  of  the  other  metals.  Many  of 
the  salts  of  lithium  and  of  sodium  form  stable  hydrates ;  the  salts  of 
potassium,  rubidium,  and  caesium,  on  the  other  hand,  are  nearly  all 
anhydrous. 


T 


CHAPTER  XXVII 
ELEMENTS   OF   GROUP   I.—  SUB-GROUP   B 

HIS  sub-group  includes  the   following   three   metals,  generally 
known  as  the  members  of  the  copper  group  :  — 

Copper  (Cu)          .....       63.57 

Silver  (Ag)    ......     107.88 

Gold(Au)      ......     197.2 

These  elements  show  only  a  comparatively  distant  resemblance  to 
the  alkalis,  although  both  belong  to  the  first  group.  They  are  all 
univalent  elements,  but  whilst  silver  only  functions  with  this  valency, 
copper  also  acts  as  a  divalent  and  gold  as  a  trivalent  element.  The 
metals  differ  markedly  from  the  alkalis  in  having  small  affinity  for 
oxygen  ;  they  also  melt  at  high  temperatures. 

COPPER 

Symbol,  Cu.     Atomic  weight,  63.57. 

History  —  On  account  of  its  small  affinity  for  oxygen,  and  the 
consequent  readiness  with  which  it  can  be  obtained  from  certain  ores 
(e.g.  the  carbonate),  copper  has  been  known  from  the  earliest  times. 
According  to  Berthelot,  a  copper  age  followed  the  stone  age  and 
preceded  the  bronze  age.  Utensils  consisting  almost  entirely  of 
copper  were  in  use  in  Egypt  and  Babylonia  anterior  to  4500  B.C. 

Occurrence  —  Copper  occurs  free  in  various  parts  of  the  world  ; 
in  great  quantity  near  Lake  Superior  and  in  New  Mexico  (United 
States).  In  the  combined  state  it  is  very  widely  distributed  ;  the 
more  important  ores  are  as  follows  :  — 

Ruby  ore  (cuprite),  Cu2O.  Copper  pyrites  or  chalcopyrite, 
Malachite,  CuCOs,Cu(OH)2.  Cu£,Fe£s,  or  CuFeS2. 

Azurite,  sCuCO3,Cu(OH)2.  Purple  copper  ore,  3Cu2S,Fe2S3> 
Copper  glance  or  chalcocite,  GuyS.  or  Cu3FeS3. 


Metallurgy  of  Copper  —  (i)  From  N  on-  sulphur  Ores  —  The 
preparation  of  copper  from  the  oxide  and  carbonate  is  very  simple, 

423 


424    A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 


the  ores  being  mixed   with   coal  or  coke  and  smelted  in  a  blast- 
furnace : 


Naturally  occurring  copper  is  separated  from  the  accompanying 
impurities  by  grinding  and  washing.  The  metal  obtained  by  these 
or  other  processes  can  be  purified  by  the  electrolytic  method  (see 
below). 

(2)  From  Sulphide  Ores  —  These  ores  usually  contain  iron  as  well 
as  sulphur,  and  their  removal  presents  considerable  difficulty,  mainly 
owing  to  the  fact  that  sulphur  has  a  greater  affinity  for  copper  than 
for  iron,  and,  further,  copper  sulphide,  unlike  the  sulphides  of  iron 
and  certain  other  metals,  is  only  converted  with  great  difficulty  into 


FIG.  85. 

the  oxide  by  roasting  in  air.  The  first  step  consists  in  roasting  the 
ore  in  a  reverberatory  furnace  (Fig.  85),  whereby  volatile  impurities 
(arsenic,  antimony,  etc.),  are  burned  off,  and  the  iron  sulphide  is 
partially  converted  to  oxide.  The  roasted  ore  is  then  heated  to 
fusion  (smelted),  whereby  a  mixture  of  cuprous  and  iron  sulphides, 
known  as  "matte"  or  "coarse  metal,"  is  obtained.  The  matte  is 
then  fused  with  sand,  whereby  a  readily  fusible  slag  of  iron  silicate 
is  formed,  which  floats  on  the  copper  and  can  be  run  off.  This 
alternate  roasting  and  fusing  is  repeated  several  times  till  all  the  iron 
is  removed  and  a  mixture  of  copper  and  cuprous  sulphide  remains. 
This  is  then  heated  in  a  reverberatory  furnace,  when  interaction 
between  sulphide  and  oxide  occurs  : 

Cu2S  +  2Cu2O->6Cu  +  SO2. 
The  impure  copper  thus  obtained  must  now  be  refined.     This  may  be 


COPPER  425 

done  in  a  furnace  or  electrolytically.  According  to  the  former  method 
it  is  melted  in  a  reverberatory  furnace  in  a  current  of  air,  whereby 
the  impurities  are  oxidized,  and  either  volatilize  or  combine  with  the 
silica  lining  the  hearth  to  form  a  slag.  Finally,  the  oxide  of  copper, 
formed  to  some  extent  in  this  process,  is  reduced  by  stirring  the 
melted  mass  with  poles  of  fresh  wood ;  the  gases  formed  by  the 
burning  of  the  wood  effect  the  reduction  of  most  of  the  cuprous  oxide. 

Instead  of  the  lengthy  process  above  indicated,  sulphide  ores  are 
now  worked  up  by  the  converter  method.  The  ore  is  first  roasted, 
then  fused,  and  placed  in  a  Bessemer  converter  (p.  555)  lined  with 
silica.  A  blast  of  air  is  then  forced  through  the  mass,  whereby  the 
sulphur  is  burned  off  as  sulphur  dioxide,  the  arsenic  and  antimony 
also  escape  as  oxides,  and  the  ferrous  oxide  combines  with  the  silica 
to  form  a  slag.  The  copper  thus  obtained  is  then  refined  by  the 
furnace  method,  or,  better,  by  electrolysis. 

(3)  Wet  Methods — These  methods  of  extraction  are  used  for  ores 
which  cannot  profitably  be  treated  in  the  dry  way.  When  the  ore  is 
free  from  sulphur  it  is  extracted  with  sulphuric  acid,  and  the  solution, 
which  contains  the  copper  as  cupric  sulphate,  treated  with  ferrous 
chloride  and  sulphur  dioxide,  whereby  reduction  to  cuprous  chloride 
takes  place.  From  the  latter  the  copper  is  obtained  by  displacement 
with  iron  : 

2CuCl  +  Fe->2Cu  +  FeCl2. 

A  method  in  use  for  sulphide  ores  is  to  roast  with  common  salt, 
whereby  cupric  chloride  is  formed  ;  the  latter  is  then  extracted  with 
water,  and  the  copper  precipitated  by  addition  of  scrap  iron  : 

CuCl2  +  Fe-»FeCl2  +  Cu. 

Electrolytic  Refining:  of  Copper — Since  impurities  such  as 
iron,  cuprous  oxide,  etc.,  greatly  diminish  the  value  of  copper  for 
commercial  purposes,  it  is  important  to  have  a  convenient  method 
for  obtaining  the  pure  metal,  and  the  requirements  are  fully  met  by 
the  electrolytic  method.  Thick  plates  of  fairly  pure  copper,  which 
form  the  anode,  are  suspended  in  an  electrolyte  containing  8  to  14 
per  cent,  of  copper  sulphate  and  4  to  10  per  cent  of  sulphuric  acid. 
Thin  plates  of  copper  form  the  cathode.  When  the  current  passes, 
the  copper  is  dissolved  from  the  anode  and  deposited  in  a  very  pure 
state  on  the  cathode,  so  that  the  net  result  of  the  process  is  the  con- 
veyance of  copper  from  one  pole  to  the  other.  Some  of  the  impurities, 


426     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

such  as  silver,  antimony,  and  cuprous  oxide,  do  not  dissolve,  and  are 
found  in  the  so-called  anode  mud  ;  other  impurities,  such  as  zinc,  dis- 
solve at  the  anode  and  remain  in  solution.  Finally,  the  pure  copper 
is  stripped  from  the  cathodes,  melted,  and  cast  in  blocks. 

As  there  is  very  little  polarization  in  this  process,  a  small  E.M.F. 
(less  than  i  volt)  is  sufficient. 

Properties  —  Copper  is  a  bright-red  metal  of  density  8.94  to 
8.96  ;  it  melts  at  1084°,  and  boils  at  2310°,  in  absence  of  air.  It  is 
fairly  hard,  but  very  tough  and  flexible  ;  it  can  be  drawn  out  into 
thin  wire  and  beaten  into  thin  sheets.  It  is  a  very  good  conductor 
of  electricity,  being  surpassed  in  this  respect  only  by  silver.  Its 
conductivity  and  malleability  are  greatly  diminished  by  traces  of 
impurities. 

In  dry  air  copper  becomes  coated  with  a  very  thin  film  of  oxides, 
which  scarcely  affect  the  colour,  but  protect  it  against  further  oxida- 
tion. In  moist  air  it  becomes  covered  with  a  thin  coating  of  basic 
carbonate  (malachite).  At  a  red  heat  it  combines  fairly  rapidly  with 
oxygen  to  form  black  cupric  oxide. 

Copper  is  not  acted  upon  by  water  at  any  temperature,  nor  is  it 
affected  by  dilute  acids  (other  than  oxyacids)  at  room  temperature  in 
absence  of  air.  Nitric  acid  rapidly  dissolves  copper  at  the  ordinary 
temperature  (p.  225),  cupric  nitrate  being  formed.  When  heated 
with  concentrated  sulphuric  acid,  cupric  sulphate  is  formed  and  sul- 
phur dioxide  given  off.  Finely-divided  copper  dissolves  slowly  on 
boiling  with  concentrated  hydrochloric  acid,  with  formation  of  cuprous 
chloride  : 


With  free  access  of  air,  copper  dissolves  slowly  even  in  dilute  acids, 
as  a  result  of  the  simultaneous  action  of  the  latter  and  oxygen,  e.g. 


Under    the    same    circumstances,    copper   is    rapidly    dissolved    by 
ammonia,  with  formation  of  a  deep  blue  solution  : 

Cu  +  O  +  H20  +  2NH3-»Cu(NH3)2(OH)2. 

On  account  of  its  stability  in  the  air  and  resistance  to  acids,  copper 
is  largely  used  commercially  for  coinage,  for  covering  the  hulls  of 
ships,  for  conveying  electricity,  etc.,  and  also  in  electrotyping.  Some 
important  alloys  of  copper  are  mentioned  in  the  next  section. 


COPPER  427 

Alloys  Of  Copper — The  more  important  are  brass,  which  con- 
tains 16  to  35  per  cent,  of  zinc  ;  the  bronzes,  which  consist  of  copper, 
zinc  and  tin  in  varying  proportions,  and  usually  lead  as  well ;  German 
silver,  which  contains  2  parts  of  copper,  i  part  of  nickel,  and  i  part 
of  zinc,  and  is  nearly  white  ;  gun-metal,  copper  with  10  per  cent,  tin  ; 
bell-metal,  copper  with  20  to  25  per  cent,  tin  ;  aluminium  bronze, 
copper  with  5  to  10  per  cent,  of  aluminium  ;  phosphor  bronze,  copper 
with  5  to  15  per  cent,  of  tin  and  0.25  to  2.5  per  cent,  of  phosphorus  ; 
manganese  bronze,  copper  with  30  per  cent,  of  manganese.  The 
copper  coinage  of  this  country  contains  95  parts  of  copper,  4  parts 
of  tin  and  i  part  of  zinc. 

Compounds  of  Copper — Copper  is  the  first  metal  we  have 
met  with  which  forms  more  than  one  series  of  salts.  In  cuprous 
salts,  which  are  of  the  type  CuX  (X  =  univalent  anion),  the  copper  is 
univalent,  whilst  in  cupric  salts,  of  the  type  CuX2,  the  copper  is 
divalent.  Cuprous  salts  of  oxygen  acids  are  practically  unknown, 
but  the  cupric  salts  of  these  acids  are  quite  stable.  The  relative 
stability  of  cuprous  and  cupric  salts  depends  greatly  on  the  condi- 
tions. At  high  temperatures  the  cuprous  halid£s_are_the_mQre  stable, 
puprous  salts  are  usually  prepared  by  reducing  cupric  salts,  and  they 
tend  to  absorb  oxygen  from  the  air,  with  reformation  of  derivatives 
of  divalent  copper.  Pure  cuprous  salts  with  a  colourless  anion  are 
colourless  (the  Cu'  ion  appears  to  be  colourless) ;  cupric  salts  in  dilute 
solution  are  greenish-blue  (that  is,  the  Cu-  ion  is  blue). 


CUPROUS  SALTS 

Cuprous  Oxide,  Cu2O — This  substance  occurs  naturally  as 
ruby  copper  ore.  It  is  obtained,  mixed  with  cupric  oxide,  by  gently 
heating  copper  in  the  air,  and  is.  also  formed  by  the  action  of  alkalis, 
e.g.  sodium  hydroxide,  on  cuprous  chloride.  It  is  most  conveniently 
obtained  by  heating  a  cupric  salt  in  alkaline  solution  with  a  reducing 
agent  such  as  grape  sugar. 

Cuprous  oxide  is  a  bright  red  powder,  insoluble  in  water.  When 
treated  with  halogen  acids,  cuprous  halides  are  obtained  ;  with  oxy- 
acids,  on  the  other  hand,  the  corresponding  cupric  salts  are  formed 
and  metallic  copper  separates  : 

Cu2O  +  H2SO4->CuSO4  +  Cu  4-  H2O. 

It  is  soluble  in  ammonia,  and  the  solution,  which   is  colourless  in 
absence  of  air,  contains  the  compound,  Cu(NH3)OH. 


428     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Cuprous  Chloride,1  CuCl  —  This  substance  is  obtained  by  dis- 
solving cuprous  oxide  in  hydrochloric  acid,  or,  more  conveniently,  by 
boiling  cupric  chloride  with  hydrochloric  acid  and  copper  turnings, 
and  then  pouring  the  solution  into  boiled  water  : 

CuCl2+Cu->2CuCl. 

Cuprous  chloride  settles  out  as  a  white  precipitate.  When  rapidly 
washed  with  glacial  acetic  acid,  then  with  absolute  alcohol  and 
quickly  dried,  it  is  stable  in  dry  air.  In  moist  air  it  turns  green,  due 
to  the  formation  of  cupric  oxychloride,  Cu(OH)Cl  or  CuCl2*Cu(OH)2- 
When  treated  with  hot  water,  cuprous  oxide  is  formed  by  hydrolysis. 
With  cold  water,  the  hydrolytic  action  is  accompanied  by  partial 
decomposition  according  to  the  equation 

2CuCl->CuCl2  +  Cu. 

Cuprous  chloride  is  soluble  in  hydrochloric  acid  with  formation  of 
the  complex  compound,  HCuCl2,  and  in  ammonia,  with  formation  of 
an  addition  compound,  Cu(NH3)2Cl.  Both  solutions  are  colourless 
when  fresh,  but  in  contact  with  air  they  rapidly  become  blue  owing 
to  the  formation  of  cupric  salts  by  oxidation.  The  solutions  have  the$ 
power  of  rapidly  absorbing  carbon  monoxide,  and  are  used  for  this 
purpose  in  gas  analysis.  The  absorption  appears  to  depend  mainly 
upon  the  formation  of  an  unstable  compound,  2CuCl,CO,2H2O,  which 
has  been  isolated  in  the  solid  form. 

Cuprous  Iodide,  Cul,  is  obtained  as  a  precipitate  when  potas- 
sium iodide  is  added  to  a  solution  of  a  cupric  salt  : 

2CuSO4+4KI->2K2SO4  +  2CuI  +  12. 

We  may  assume  that  cupric  iodide,  CuI2,  is  first  formed,  but,  being 
unstable,  it  decomposes  immediately  into  cuprous  iodide  and  iodine. 

Cuprous  iodide  occurs  in  small,  colourless,  octahedral  crystals, 
practically  insoluble  in  water. 

Cuprous  Cyanide,  CuCN,  is  obtained,  analogous  to  the  iodide, 
by  warming  a  solution  of  cupric  sulphate  with  potassium  cyanide  : 


Cyanogen  escapes  as  a  gas,  and  cuprous  cyanide  is  obtained  as  a 

1  Cuprous  salts  are  sometimes  written  with  the  double  formula,  cuprous 
chloride,  for  example,  as  Cu2Cl2,  but  molecular  weight  determinations  in  solution 
and  the  way  in  which  the  compounds  react  (cf.  p.  132)  lend  support  to  the 
simple  formulae. 


COPPER  429 

white  precipitate,  very  slightly  soluble  in  water,  but  readily  soluble 
in  a  solution  of  potassium  cyanide.  The  latter  solution  contains 
a  compound  KCu(CN)2  which  is  partly  dissociated  into  K'  and 
Cu(CN)2/  ions. 

Cuprous  Sulphide,  Cu2S,  occurs  naturally  as  copper  glance, 
and  can  be  prepared  by  direct  combination  of  the  elements  or  by 
heating  cupric  sulphide  in  a  current  of  hydrogen  at  260°.  It  is  a 
dark  brown  crystalline  powder. 

CUPRIC  SALTS 

Cupric  Oxide,  CuO — This  compound  is  obtained  by  strongly 
heating  copper  in  air  or  oxygen  or  by  heating  cupric  hydroxide, 
carbonate  or  nitrate.  It  is  a  black  powder,  and  has  considerable 
power  for  absorbing  gases,  such  as  carbon  dioxide  and  water  vapour. 
When  heated  to  1000°,  it  loses  part  of  its  oxygen,  and  a  mixture  of 
cuprous  and  cupric  oxides  is  obtained.  It  gives  up  its  oxygen  fairly 
readily  to  reducing  substances,  such  as  free  hydrogen,  organic  com- 
pounds, etc.,  and  for  this  reason  is  largely  used  in  organic  analysis. 

Cupric  Hydroxide,  Cu(OH)2,  is  formed  as  a  blue  flocculent 
precipitate  when  sodium  or  potassium  hydroxide  is  added  to  the 
solution  of  a  cupric  salt.  When  the  liquid  in  which  it  is  suspended 
is  boiled  it  loses  water,  black  cupric  oxide  being  formed. 

Cupric  hydroxide  readily  dissolves  in  ammonia,  and  the  resulting 
solution,  which  probably  contains  the  compound  Cu(NH3)4(OH)2,  has 
the  remarkable  property  of  dissolving  cellulose. 

Cupric  hydroxide  is  a  very  weak  base,  and  therefore  all  the  cupric 
salts  have  an  acid  reaction  owing  to  hydrolysis. 

Cupric  Chloride,  CuCl2,  is  obtained  by  dissolving  cupric  oxide 
or  carbonate  in  hydrochloric  acid,  and  separates  in  blue  needles  of 
the  dihydrate,  CuCl2,2H2O,  on  concentrating  the  solution.  The  anhy- 
drous salt  can  be  prepared  by  heating  the  dihydrate  in  a  current  of 
dry  hydrogen  chloride,  and  is  brownish-yellow. 

The  concentrated  aqueous  solution  of  cupric  chloride  is  green,  but 
on  progressive  dilution  the  colour  gradually  changes  and  finally  attains 
the  deep  blue  colour  characteristic  of  dilute  solutions  of  all  cupric 
salts.  The  simplest  explanation  of  these  observations  is  that  con- 
centrated solutions  contain  almost  exclusively  the  non-ionised  salt, 
and  that  as  the  solution  is  diluted  the  colour  approximates  more  and 
more  to  that  of  the  divalent  Cu"  ion  (or,  more  probably,  of  the 
hydrated  ion  Cu(H2O)4").  Although  the  increased  ionisation,  accom- 
panied by  increased  hydration,  is  no  doubt  the  main,  it  is  not  the 


430     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

only  factor  concerned  in  the  changes  of  colour,  as  it  has  been  shown 
that  in  concentrated  solutions  complex  cations  containing  chlorine, 
e.g.  Cud',  are  also  present. 

When  ammonia  is  added  in  excessjo  a  solution  of  cupric  chloride 
a  deep  azure-blue  solution  is  'obTained,  from  which  the  compound 
Cu(NH3)4Cl2,H2O  separates  on  concentration  in  dark  blue  needles. 
The  compound  Cu(NH3)6Cl2,  a  dark  blue  powder,  is  obtained  by  the 
action  of  ammonia  gas  on  the  anhydrous  salt.  When  these  com- 
pounds are  heated  to  150°  ammonia  is  given  off,  and  the  compound 
CuCl2,2NH3  remains  as  a  dark  green  powder. 

Compounds  of  Copper  Salts  with  Ammonia— When 
ammonium  hydroxide  is  added  to  any  cupric  salt,  azure-blue  solutions 
are  obtained  which  contain  compounds  of  the  salts  with  ammonia. 
There  is  evidence  that  the  copper  in  these  solutions  (including  that 
of  cupric  hydroxide)  is  present  mainly  if  not  entirely  in  the  form 
of  salts  of  the  type  Cu(NH3)4X2,  which  on  dissociation  give  rise 
to  divalent  Cu(NH3)4"  ions.  Such  salts  can  also  be  isolated  in  the 
solid  form,  but  the  composition  in  this  case  is  very  varied. 

Cuprous  salts  also  dissolve  in  ammonia  giving  colourless  solutions 
which  very  rapidly  become  blue  in  the  air  owing  to  the  formation 
of  cupric  compounds  by  oxidation.  In  some  cases  at  least  these 
solutions  contain  complex  salts  of  the  type  Cu(NH3)2X,  which  yield 
cations  of  the  type  Cu(NH3)2'  by  ionisation. 

Cupric  Sulphate  (blue  vitriol}  CuSO4 — This  salt  is  obtained 
by  heating  copper  with  concentrated  sulphuric  acid,  or  by  dissolving 
the  oxide  or  carbonate  in  dilute  acid.  On  the  commercial  scale,  it 
is  prepared  by  roasting  copper  pyrites  in  a  current  of  air,  extracting 
with  water  and  evaporating,  when  the  salt  separates  in  large  blue 
triclinic  crystals  of  the  formula  CuSO4,5H2O. 

On  heating  for  some  time  at  1 10°  the  salt  loses  4.H2O,  and  becomes 
anhydrous  when  heated  for  some  time  at  210°.  The  anhydrous  salt 
is  grayish-white,  has  a  great  affinity  for  water,  and  is  used  as  a 
dehydrating  agent.  On  heating  more  Strongly,  sulphur  dioxide  and 
oxygen  are  given  off,  and  basic  salts  of  the  type  *CuSO4lyCuO  are 
formed. 

At  15°,  30°,  and  50°,  the  solubility  in  100  grams  of  water  is  19.3, 
25.5,  and  33.6  grams  respectively,  calculated  on  the  anhydrous 
salt.  , 

The  behaviour  of  copper  sulphate  towards  ammonia  is  similar  to 
that  of  the  chloride.  From  the  azure-blue  solution  containing  excess 
of  ammonia  the  compound  Cu(NH3)4SO4,H2O  has  been  obtained 


COPPER  43i 

in  blue  crystals.  A  number  of  other  solid  compounds  have  been 
described,  but  their  existence  is  somewhat  doubtful. 

Cupric  Nitrate,  Cu(NO3)2,  is  obtained  by  dissolving  copper, 
the  dioxide  or  carbonate  in  dilute  nitric  acid  and  evaporating  the 
solution.  Above  24.5°  a  trihydrate  separates  from  solution,  at  lower 
temperatures  a  hexahydrate,  in  blue  tabular  crystals,  and  at  still 
lower  temperatures  the  compound  Cu(NO3)2,9H2O. 

Cupric  Carbonates — The  normal  carbonate  of  copper  is  not 
known.  When  a  solution  of  copper  sulphate  is  heated  with  sodium 
carbonate,  the  bluish-green  colloidal  precipitate  (of  varying  com- 
position) first  obtained  changes  on  standing  in  contact  with  the  mother 
liquor  to  the  crystalline  basic  carbonate  3CuCO3,3Cu(OH)2,H2O. 
Important  naturally  occurring  basic  carbonates  are  malachite, 
CuCO3,Cu(OH)2,  and  azurite,  2CuCO3,Cu(OH)2,  both  of  which  have 
been  prepared  artificially. 

Cupric  Sulphide,  CuS,  is  obtained  as  a  black  precipitate  by 
passing  hydrogen  sulphide  through  a  solution  of  a  cupric  salt.  It  is 
practically  insoluble  in  water  and  in  hydrochloric  acid,  but  is  slightly 
soluble  in  yellow  ammonium  sulphide. 

Tests  for  Copper — Copper  salts  (in  the  cupric  form)  are 
characterized  by  the  blue  colour  of  their  solutions,  by  the  azure-blue 
solution  obtained  on  adding  excess  of  ammonia,  and  by  the  formation 
of  a  black  precipitate  of  cupric  sulphide  when  hydrogen  sulphide  is 
passed  through  a  solution. 

Oxidation  and  Reduction— Up  to  the  present  we  have  re- 
garded the  process  of  oxidation  as  being  associated  with  the  addition 
of  oxygen  to,  or  the  removing  of  hydrogen  from,  a  compound ;  con- 
versely, when  oxygen  is  removed  from,  or  hydrogen  added  to,  a 
compound  the  latter  is  said  to  be  reduced.  Free  oxygen,  many 
oxides  and  peroxides,  e.g.  manganese  peroxide,  and  many  compounds 
rich  in  oxygen,  e.g.  nitric  acid,  chloric  acid,  are  familiar  oxidizing 
agents.  On  the  other  hand,  free  hydrogen,  "nascent"  hydrogen 
(p.  514)  and  hydriodic  acid  are  reducing  agents. 

Many  substances  which  contain  no  oxygen  act  as  oxidizing  agents, 
e.g.  chlorine,  and  similarly,  substances  containing  no  hydrogen  may 
act  as  reducing  agents.  These  indirect  oxidizing  and  reducing  agents 
usually  act  through  the  intervention  of  water.  Thus  chlorine  combines 
with  water  to  form  hydrogen  chloride,  the  oxygen  becoming  available 
for  oxidizing  purposes  (p.  86) : 

2C12-|-2H2O->4HC1  +  O2, 


432     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

and  similarly,  sulphur  dioxide,  when  it  acts  as  a  reducing  agent,  gives 
sulphuric  acid  and  hydrogen  : 

S02+H20->S03+H2 
SO3  +  H2O-»H2SO4. 

It  has  been  repeatedly  pointed  out  that  oxidation  is  always  accom- 
panied by  reduction  ;  while  one  substance  is  oxidized,  the  oxidizing 
agent  undergoes  reduction. 

In  connexion  with  metals  such  as  copper,  which  have  two  valencies, 
it  is  necessary  to  extend  somewhat  the  use  of  the  terms  oxidation  and 
reduction.  As  cuprous  salts  are  converted  to  cupric  salts  by  the 
action  of  oxidizing  agents,  and  conversely,  cupric  salts  are  converted 
to  cuprous  salts  by  reducing  agents,  we  say  that  cupric  chloride, 
CuCl2,  for  instance,  represents  a  higher  state  of  oxidation  than  cuprous 
chloride,  CuCl,  although  neither  oxygen  nor  hydrogen  are  directly 
concerned.  When  we  consider  that  the  oxide  corresponding  with 
cuprous  chloride  is  Cu2O,  and  that  corresponding  with  cupric  chloride 
is  CuO,  it  is  evident  that  this  way  of  representing  the  matter  is 
justified. 

The  same  considerations  apply  to  metals  as  compared  with  salts. 
As  the  oxide  corresponding  with  cuprous  chloride  is  Cu2O,  it  is  usual 
to  speak  of  the  conversion  of  cuprous  chloride  (or  any  other  salt  of 
copper)  to  the  metal  as  reduction,  and  the  formation  of  the  salt  from 
the  metal  as  oxidation.  It  follows  that  the  combination  of  a  metal 
with  almost  any  element  (except  hydrogen  or  another  metal)  may  be 
regarded  as  oxidation. 

The  matter  becomes  clearer  on  the  basis  of  the  electrolytic  dis- 
sociation theory.  Comparing  copper,  cuprous  chloride  and  cupric 
chloride,  we  may  say  that,  as  far  as  salts  are  concerned,  an  increase 
in  the  number  of  positive  charges  (or  a  diminution  in  the  number 
of  negative  charges)  on  an  ion  denotes  oxidation;  decrease  in  the 
number  of  positive  charges  or  an  increase  in  the  number  of  negative 
charges  on  an  ion  denotes  reduction.  Further  illustrations  of  these 
statements  will  be  given  later. 

Electromotive  Force1  — At  a  very  early  stage  of  our  work, 
we  saw  that  when  zinc  and  copper  are  connected  by  a  wire  and 
dipped  into  dilute  sulphuric  acid,  an  electric  current  passes  through 
the  wire  (p.  13).  This  process  was  regarded  as  a  transformation  of 
chemical  into  electrical  energy,  the  source  of  the  energy  being  the 

i  For  full  details,  see  Physical  Chemistry,  chap.  xiv. 


COPPER  433 

oxidation  of  the  zinc,  which  in  the  process  is  transformed  into  zinc 
sulphate.  The  hydrogen  is  liberated  at  the  surface  of  the  copper. 

An  arrangement  of  this  sort  is  termed  a  galvanic  cell  or  battery. 
For  our  present  purpose,  it  is  rather  more  instructive  to  use  a 
solution  of  copper  sulphate  as  electrolyte  in  contact  with  the  copper 
instead  of  the  sulphuric  acid  previously  used.  Further,  although 
a  current  is  obtained  when  the  electrodes  are  simply  placed  at 
some  distance  apart  in  the  electrolyte,  it  is  preferable,  in  order  to 
minimize  mixing  of  the  solutions  round  the  two  poles,  to  separate 
them  by  some  arrangement  which  at  the  same  time  allows  the 
current  to  pass.  An  earthenware  pot  answers  the  purpose  satis- 
factorily. The  battery  we  are  considering,  then,  is  made  by  placing 
in  the  outer  vessel  a  solution  of  copper  sulphate  in  which  a  copper 
plate  is  partially  immersed,  and  in  the  inner  vessel  —  a  porous 
earthenware  pot— a  rod  of  zinc  dipping  in  a  solution  of  sulphuric 
acid.  When  the  poles  are  connected  by  a  wire,  an  electric  current 
passes  round  the  circuit — in  the  solution  positive  electricity  flows 
from  zinc  to  copper,  and  in  the  connecting  wire  from  copper  to 
zinc — and  simultaneously  zinc  dissolves  at  the  zinc  pole  and  an 
equivalent  of  copper  is  deposited  on  the  copper  pole.  The 
direction  of  the  current  is  shown  in  the  accompanying  diagram. 

In  the  external  circuit. 
Zn  |  H2SO4  |  CuSO4  |  Cu 

In  the  solution. 

The  pole  where  positive  electricity  leaves  the  solution — in  this  case 
the  copper  plate — is  termed  the  positive  pole,  the  electrode  at  which 
positive  electricity  enters  the  solution  is  called  the  negative  pole. 

The  chemical  change  accompanying  the  passage  of  the  current 
may  be  written  thus  : 

Zn  +  CuSO4->ZnSO4+  Cu, 
or,  from  the  ionic  point  of  view  : 

Zn  +  Cu '  '->Zn  "  -f  Cu  ; 

that  is,  it  is  simply  the  displacement  of  copper  from  solution  by 
zinc.  This  change,  as  we  have  already  seen,  takes  place  directly  ; 
when  zinc  is  put  in  a  solution  of  copper  sulphate,  copper  is  deposited 
and  an  equivalent  of  zinc  goes  into  solution.  As  the  change  takes 
28 


434    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

place  spontaneously,  we  anticipate  that  there  is  a  diminution  of 
energy  in  the  process,  and  this  energy  is  obtained  in  the  galvanic 
cell  as  electrical  energy. 

A  word  is  necessary  regarding  the  factors  of  electrical  (energy. 
The  work  which  an  electric  current  can  do  depends  not  only  on 
the  quantity  of  electricity  which  passes,  but  on  the  potential  or 
electromotive  force  in  the  circuit.  The  quantity  of  electricity  is 
measured  in  coulombs,  the  electromotive  force  in  volts.  '  Electrical 
energy  is  proportional  to  the  product  of  the  electromotive  force 
and  the  quantity  of  the  electricity.  Thus  the  same  amount  of 
light  will  be  given  out  when  5  coulombs  pass  through  an  electric 
glow-lamp  at  a  potential  of  10  volts,  as  when  2  coulombs  pass 
through  at  a  potential  of  25  volts. 

The  Potential  Series  of  the  Elements— It  can  be 
shown  that  the  electromotive  force  of  the  zinc-copper  cell,  arranged 
as  we  have  described,  is  rather  more  than  one  volt.  If  magnesium 
is  used  in  place  of  zinc,  the  E.M.F.  is  about  1.8  volts.  Similar  cells 
can  be  built  up  with  other  metals  in  place  of  zinc  and  copper,  and 
in  every  case,  when  one  metal  can  displace  the  other  from  combina- 
tion an  electric  current  is  obtained.  Now  it  can  be  shown  that  the 
E.M.F.  of  a  cell  built  up  with  two  metals  in  this  way  is  a  measure 
of  the  tendency  of  one  element  to  displace  the  other  from  combina- 
tion. Proceeding  on  these  lines,  the  metals  have  been  arranged  in 
order  in  such  a  way  that  each  metal  can  displace  from  combination 
all  those  following  it  in  the  series,  and  is  displaced  from  combination 
by  all  those  preceding  it  in  the  series. 

The  so-called  potential  series  of  the  metals  is  as  follows,  the 
potential  of  hydrogen  being  taken  as  zero : 

Na         Mg          Al          Mn          Zn          Cd  Fe          Co 

-2.58  —1.482  —1.276  —1.075    -0.770   -0.420  -0.334  —0.232 

Ni  Pb         H2         Cu          Hg        Ag        Pt         Au 

—0.228     -0.151     ±o     +0.329    0.753    0.771     1.42?     1.65? 

These  numbers  represent  the  potential  difference  between  a  metal 
and  the  solution  of  one  of  its  salts  in  normal  solution,  referred  to 
the  potential  between  a  normal  solution  of  an  acid  and  a  platinum 
electrode  saturated  with  hydrogen,  which  is  taken  as  zero.  The 
actual  zero  chosen  is  a  matter  of  indifference;  the  important  point 
is  the  difference  of  potential  between  one  substance  and  another. 
The  table  shows  that  magnesium,  aluminium  and  zinc  can  displace 


SILVER 


435 


iron  from  its  salts  ;  iron,  on  the  other  hand,  can  displace  lead,  copper 
and  silver  from  combination. 

A  very  interesting  point  is  the  light  thrown  by  this  list  on  the 
behaviour  of  the  metals  with  regard  to  the  displacement  of  hydrogen 
from  acids.  All  the  metals  preceding  hydrogen  in  the  list  should 
be  able  to  liberate  it  from  acids,  and  we  have  seen  in  a  number 
of  instances  (others  are  given  later)  that  such  is  the  case.  On  the 
other  hand,  copper,  mercury  and  silver  cannot  displace  hydrogen 
from  acids  ;  on  the  contrary,  hydrogen  should  be  able  to  displace 
these  metals  from  their  salts.  This  is  not  observed  under  ordinary 
circumstances,  probably  owing  to  the  reactions  being  exceedingly 
slow,  but  hydrogen  occluded  in  platinum  can  displace  some  of  the 
metals  from  combination  owing  to  the  accelerating  influence  of 
platinum  on  the  rate  of  reaction. 

Another  method  of  stating  the  facts  is  that  when  the  metals  are 
arranged  in  the  order  of  their  potentials,  they  are  arranged  in  the 
order  of  their  solution  pressures  (p.  85).  Metals  such  as  sodium 
have  so  high  a  solution  pressure  that  they  decompose  water  at  room 
temperature ;  the  solution  pressures  of  copper  and  mercury,  on 
the  other  hand,  are  extremely  small.  The  metal,  of  course,  goes  into 
solution  as  positive  ions,  this  being  the  only  form  in  which  it  can 
dissolve  in  ordinary  solvents. 

An  interesting  application  of  these  principles  is  met  with  in  the  case 
of  iron  coated  with  zinc  (p.  466)  and  tin  (p.  496)  respectively.  If  such 
coatings  are  broken  at  any  point  an  electrolytic  cell  is  set  up  in  the 
presence  of  water  containing  carbon  dioxide.  Owing  to  its  position 
in  the  potential  series  zinc  dissolves  and  hydrogen  is  given  off  at  the 
iron  which  is  not  affected,  but  in  the  presence  of  tin  the  iron  dissolves 
(cf.  table)  and  rusting  is  thereby  facilitated. 

SILVER 

Symbol,  Ag.     Atomic  Weight  — 107.88. 

Chemical  Relations  —  In  its  compounds  silver  is  almost 
invariably  univalent.  The  hydroxide  is  a  fairly  strong  base  and 
therefore  silver  salts  are  not  appreciably  hydrolyzed.  Like  copper, 
it  has  a  tendency  to  form  complex  ions.  It  has  very  little  affinity 
for  oxygen,  and  silver  salts  are  therefore  readily  reduced  to  the 
metallic  state. 

Occurrence— Silver  occurs  free  in  nature,  usually  in  small  pro- 
portion distributed  in  quartz  and  other  rocks,  occasionally  in  large 
masses.  In  combination  it  occurs  as  sulphide  mixed  with  lead 


436     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

sulphide  (galena),  with  which  it  is  isomorphous.  The  more  im- 
portant silver  ores  are  :  silver  glance  or  argentite,  Ag2S  ;  pyrargyrite 
or  ruby  silver  ore,  3Ag2S,Sb2S3  or  Ag3SbS3;  proustite,  3Ag2S,As2S3 
or  Ag3AsS3,  and  horn  silver,  AgCl.  Silver  is  always  present  in 
native  copper  and  (as  sulphide)  in  copper  pyrites  (p.  423).  The 
supply  of  silver  comes  mainly  from  the  United  States,  Mexico, 
Bolivia,  Germany  and  Australia. 

Metallurgy  of  Silver — The  processes  used  in  obtaining 
silver  depend  upon  the  nature  of  the  ores  in  which  it  occurs. 
When  metallic  lead  is  prepared  from  galena  containing  silver 
sulphide,  the  lead  contains  all  the  silver  as  metal,  and  its  separa- 
tion is  an  important  industry.  Two  processes  are  in  use  for  this 
purpose  :  (i)  the  Pattinson  process,  (2)  the  Parkes  process. 

(1)  The  Pattinson  Process — The  alloy  of  silver  and  lead,  which 
is  relatively  poor  in  silver  (as  little  as  0.02  per  cent,  is  commercially 
workable)  is  melted  and  allowed  to  cool,  when  lead  separates  out 
first   in  crystals,  which  are   removed   (cf.  p.  197).      This  process  is 
repeated  till  the  eutectic  point  of  lead  and  silver  is  nearly  reached 
(about  2.25  per  cent,  silver).     The  alloy  is  then  subjected  to  cupella- 
tion,  which  consists  in  heating  it  in  a   reverberatory  furnace,  the 
hearth   of  which   is   lined  with   bone-ash,   while  a   blast   of  air  is 
passed  over  the  surface.      The  lead  is  thus  converted  into  oxide, 
which  at  the  high  temperature  melts  and  is  driven   by  the   blast 
of  air  over  the  edge  of  the  "cupel."      In  course  of  time  all  the 
lead  is  thus   removed,  and  its   disappearance  is   indicated  by  the 
sudden  appearance — so-called   "  flashing  " — of   the   bright    metallic 
surface  of  the  silver. 

(2)  The  Parkes   Process — This   process,   which   has   now   largely 
superseded  the  Pattinson  process,  depends  upon  the  fact  that  when 
zinc  is  added  to  the  alloy  of  lead  and  silver,  and  the  metals  are  fused 
and  thoroughly  mixed,  the  zinc  and  lead  separate  almost  completely 
into  two  liquid  layers,  and  nearly  all  the  silver  passes  into  the  upper 
zinc  layer.    The  latter  layer  solidifies  first,  and  is  removed  by  means  of 
sieves  from  the  surface  of  the  fused  lead,  heated  gently  to  allow  most 
of  the  lead  to  run  off,  distilled  to  remove  the  zinc,  and  finally  the  still 
remaining  lead  is  removed  by  cupellation. 

Other  Processes— (i)  The  so-called  Amalgamation  Process, 
in  use  in  Mexico  since  1537,  but  now  practically  displaced  by  the 
cyanide  process,  is  carried  out  as  follows :  The  ore,  chiefly  sulphide, 
is  first  reduced  to  powder ;  mercury,  common  salt,  and  magistral  (a 
mixture  of  copper  and  iron  sulphates  produced  by  roasting  copper 


SILVER  437 

pyrites)  are  then  added,  and  thorough  mixing  is  effected  by  the 
treading  of  mules.  The  main  reactions,  which  are  complete  after 
some  weeks,  are  as  follows  : 

Ag2S  +  CuCl2->2AgCl+CuS 
2AgCl  +  2Hg-»Hg2Cl2+2Ag. 

The  mixture  is  then  stirred  up  with  water,  the  amalgam,  which  con- 
tains the  silver  dissolved  in  excess  of  mercury,  is  separated,  strained 
through  filter  bags  to  remove  the  excess  of  mercury,  and  distilled. 

According  to  an  improved  amalgamation  process  in  use  in  Germany, 
the  ores  are  roasted  with  common  salt  to  form  silver  chloride,  and  the 
latter  reduced  to  the  metal  by  rotating  the  ore  in  casks  containing 
iron  plates,  the  silver  is  then  extracted  by  means  of  mercury,  and  the 
latter  distilled  off. 

(2)  The  Cyanide  Process,  now  by  far  the  most  important  wet 
process,  is  carried  out  as  follows  :  The  ore  is  crushed  and  extracted 
with  aqueous  potassium  or  sodium  cyanide  solution  (0.1-0.5  Per  cent.), 
which  reacts  with  silver  sulphide  thus  : 

Ag2S +4NaCN$2NaAg(CN)2+ Na2S. 

The  reaction  is  reversible,  but  proceeds  almost  completely  towards 
the  right  when  the  sodium  sulphide  is  removed  by  oxidation, 
effected  by  exposure  to  air  or  by  passing  air  through  the  solution. 
After  extraction  is  practically  complete,  the  mixture  is  allowed  to 
settle,  and  the  silver  precipitated  from  the  clear  solution  by  means 
of  zinc. 

Properties — Silver  is  a  white,  lustrous  metal  which  fuses  at 
964°  and  boils  at  1950° ;  it  can  be  distilled  in  the  oxyhydrogen  flame. 
Its  density  is  10.49  to  10.50.  It  is  the  best  conductor  of  heat  and 
electricity  among  the  metals.  It  is  very  malleable  and  ductile,  and 
can  be  drawn  out  into  very  thin  wires.  Fused  silver  has  the  remark- 
able property  of  dissolving  oxygen  ;  just  above  its  melting-point  one 
volume  of  silver  takes  up  20.3  volumes  of  the  gas.  As  oxygen  is  prac- 
tically insoluble  in  solid  silver,  the  gas  escapes  as  the  metal  solidifies, 
which  gives  rise  to  the  so-called  "spitting"  of  silver,  and  causes 
curious  excrescences  on  the  solidified  metal. 

Silver  is  obtained  in  the  amorphous  form  by  reducing  silver  sulphide 
with  hydrogen.  The  gray  or  black  silver  obtained  by  reducing  the 
salts  in  solution,  e.g.  with  ferrous  sulphate,  is  ordinary  silver  in  the 
form  of  minute  needles. 


438    A   TEXT-BOOK    OF   INORGANIC    CHEMISTRY 

Silver  is  obtained  in  "colloidal"  solution  by  passing  the  electric  arc 
between  silver  rods  dipping  under  water,  or  by  reducing  silver  salts 
in  solution  with  certain  reagents,  e.g.  with  ferrous  citrate.  According 
to  the  nature  of  the  reducing  agent  and  the  conditions,  the  solutions 
differ  in  colour  (red,  green,  black,  etc.),  which  is  doubtless  connected 
with  the  degree  of  fineness  of  division  of  the  colloidal  particles. 

Silver  does  not  combine  with  oxygen  even  on  heating.  It  is 
insoluble  in  hydrochloric  or  in  dilute  sulphuric  acid,  but  on  heating 
with  concentrated  sulphuric  acid  silver  sulphate  is  formed  and  sulphur 
dioxide  given  off: 


It  is  readily  dissolved  by  nitric  acid,  silver  nitrate  being  formed  : 


Alloys  —  Pure  silver  is  rather  too  soft  to  use  for  commercial  pur- 
poses, and  the  alloy  with  copper  is  always  employed.  British  silver 
coins  contain  92.5  per  cent,  of  the  metal  ;  those  of  the  United  States 
and  continental  countries  only  90  per  cent.  For  jewellery,  plate, 
etc.,  alloys  containing  75  to  95  per  cent,  of  silver  are  used. 

Oxides  of  Silver  —  Two  oxides  of  silver,  the  normal  oxide, 
Ag2O,  and  a  peroxide,  AgO,  appear  to  be  definitely  known.  The 
existence  of  a  third  oxide,  Ag4O,  is  doubtful. 

Silver  Oxide,  Ag2O,  is  obtained  as  a  black,  amorphous  powder 
by  adding  potassium  hydroxide  to  a  solution  of  silver  nitrate,  or  by 
boiling  silver  chloride  with  potassium  hydroxide.  By  these  methods 
we  would  expect  the  formation  of  silver  hydroxide,  AgOH,  but  this 
compound  is  unstable,  and  decomposes  almost  completely  into  silver 
oxide  and  water.  The  decomposition  is  not,  however,  quite  complete, 
as  the  oxide  is  slightly  soluble  in  water,  and  the  solution  is  distinctly 
alkaline,  and  therefore  contains  silver  hydroxide.  Silver  hydroxide^ 
AgOH,  although  not  so  strong  a  base  as  the  alkalis,  is  stronger  than 
ammonium  hydroxide,  so  that  silver  salts  are  not  appreciably  hydro- 
lyzed  in  solution. 

On  heating  to  250°,  silver  oxide  decomposes  rapidly  into  silver  and 
oxygen.  It  is  reduced  to  metallic  silver  by  heating  in  hydrogen  at 
100°.  It  is  readily  soluble  in  ammonia,  and  the  solution  contains  the 
complex  compound  Ag(NH3)2OH,  which  is  a  strong  base. 

Silver  Peroxide,  AgO  (or  Ag2O2),  is  obtained  as  a  black  powder 
by  the  action  of  ozone  on  silver  or  by  the  electrolysis  of  a  solution  of 


SILVER  439 

silvei  nitrate.  As  prepared  by  the  latter  method,  it  is  always  con- 
taminated with  adhering  silver  nitrate. 

Silver  Chloride,  AgCl  —  This  salt  is  formed  as  a  white  curdy 
precipitate  when  hydrochloric  acid  or  a  soluble  chloride  is  added 
to  a  solution  of  silver  nitrate.  It  melts  at  480-490°,  and  can  be 
vaporized  without  decomposition.  On  exposure  to  light,  the  originally 
white  salt  becomes  violet,  then  brown,  and  the  odour  of  chlorine  can 
be  detected.  This  change  is  due  to  the  formation  of  a  lower  chloride 
according  to  the  equation  * 

4AgCl->2Ag2Cl  +  Cl2 

(see  silver  salts  in  photography,  below). 

The  solubility  of  silver  chloride  in  water  is  very  small,  amounting 
to  1.56  x  io~5  mols  per  litre  (i  part  in  430,000  parts  of  water)  at  25°. 
It  is  readily  soluble  in  ammonia,  in  sodium  thiosulphate,  in  potassium 
cyanide,  in  concentrated  hydrochloric  acid  and  in  concentrated  solu- 
tions of  alkali  chlorides  with  formation  of  complex  ions. 

Solubility  Product.  Complex  Ions  containing-  Silver 
—  From  the  considerations  advanced  in  previous  chapters  we  know 
that  in  a  saturated  solution  of  a  salt  at  a  definite  temperature  there 
are  two  equilibria  :  (a)  that  between  the  solid  salt  and  the  non-ionised 
salt  in  the  solution  ;  (b}  that  between  the  non-ionised  salt  and  the  ions. 
In  the  case  of  silver  chloride,  they  may  be  represented  as  follows:  — 

Ag'  +  Cl'^tAgCl  (in  solution) 

M 

AgCl  (solid). 

Further,  since  the  solution  is  saturated,  the  concentration  of  the  non- 
ionised  salt  in  the  solution  must  be  constant  at  constant  temperature, 
just  as  sugar  has  a  constant  solubility  at  a  definite  temperature. 
Applying  the  law  of  mass  action  to  the  equilibrium  in  solution  (p.  166), 
we  have 

K[AgCl]=S. 


As  the  right-hand  side  of  the  above  equation  is  constant  at  constant 
temperature,  the  product  of  the  concentration  of  the  two  ions  —  the  so- 
called  solubility  product,  S,  —  is  also  constant  at  constant  temperature. 
The  great  importance  of  these  considerations  will  be  evident. 
If  the  solubility  product  in  a  saturated  solution  of  silver  chloride 
(which  amounts  to  (1.25  x  io~6)2=i.56x  io~10  mols  per  litre  at  25°) 
is  exceeded  in  any  way,  for  example,  by  adding  CY  ions,  the 


440    A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 

equilibrium  is  displaced  towards  the  right,  and  the  excess  of  non- 
ionised  silver  chloride  falls  out  of  solution  ;  this  change  proceeds 
till  the  original  concentration  of  non-ionised  silver  chloride  is 
re-established.  If,  on  the  other  hand,  silver  ions  are  removed 
in  some  way,  the  product  of  the  ions  is  no  longer  equal  to  the 
solubility  product,  silver  chloride  will  dissolve,  and  this  change 
will  proceed  until  the  solubility  product  is  re-established.  It  is,  of 
course,  evident  that  the  Ag'  and  Cl'  ions  need  not  be  present  in 
equivalent  proportions.  Equilibrium  is  established  between  ions  and 
dissolved  non-ionised  salt  when  the  solubility  product  is  attained  with 
any  ratio  between  concentrations  of  the  ions  concerned. 

Two  illustrations  of  the  importance  of  the  solubility  product  will  be 
given.  In  quantitative  analysis  it  is  often  necessary  to  precipitate  a 
salt,  e.g.  silver  chloride,  as  completely  as  possible  from  solution.  The 
saturated  solution  of  silver  chloride  contains  about  0.0023  grams  per 
litre  of  the  salt  at  25°.  If,  however,  sodium  chloride  is  added  till  the 
solution  is  say  i/ioo  molar  with  reference  to  Cl'  ions  (that  is,  the  Cl' 
concentration  is  increased  1000  times)  silver  chloride  falls  out  of 
solution  till  the  concentration  of  Ag*  ions  is  about  i/iooo  of  its  former 
value  (as  is  clear  from  the  equation),  and  the  amount  of  silver  chloride 
remaining  in  solution  is  quite  negligible. 

The  converse  change  occurs  when  ammonia  is  added  to  a  saturated 
solution  of  silver  chloride  in  equilibrium  with  the  solid  salt.  The 
Ag'  ions  combine  with  ammonia  to  form  complex  Ag(NH3)2*  ions, 
and  the  product  of  the  ionic  concentrations  [Ag']  x  [Cl'J  falls  far  below 
the  solubility  product  ;  silver  chloride  dissolves  to  re-establish  the 
equilibrium,  and  this  proceeds  till  the  solubility  product  is  again 
reached.  Similarly,  the  solubility  of  silver  chloride  in  concentrated 
hydrochloric  acid  is  accounted  for  by  the  formation  of  complex  AgCl/ 
ions  : 


which  are  only  very  slightly  dissociated  into  Ag1  and  Cl'  ions. 

The  behaviour  of  an  unsaturated  solution  containing  complex  ions 
will  clearly  depend  upon  the  stability  of  the  complex  ion.  In  a  solu- 
tion of  silver  chloride  in  ammonia  there  are  two  equilibria  : 

Ag(N  H3)2Cl^Ag(N  H3)2-  +  Cl' 
and 


but  the  Ag*  ion  concentration    is  extremely  small.      It  follows  that 
none  of  the  ordinary  reagents  will  produce  a  precipitate  in  an  ammo- 


SILVER  44I 

niacal  solution  of  silver  chloride  unless  the  solubility  product  of  the 
substance  which  can  be  formed  is  excessively  small.  Silver  sulphide 
is  practically  the  only  silver  compound  which  answers  these  require- 
ments (solubility  product,  3.9  x  io~*°  mols  per  litre  at  25°),  and, 
therefore,  hydrogen  sulphide  causes  .a  precipitate  in  an  ammoniacal 
solution  of  silver  chloride. 

Silver  Bromide,  AgBr,  and  Silver  Iodide,  Agl,  are  prepared 
by  methods  analogous  to  those  described  for  the  chloride.  Both  are 
yellow  salts,  which  are  still  less  soluble  in  water  than  the  chloride. 
The  solubility  of  the  bromide  amounts  to  0.00012  grams,  that  of  the 
iodide  to  0.0000023  grams  per  litre  at  25°.  The  bromide  is  slightly 
soluble,  the  iodide  practically  insoluble  in  ammonia  ;  the  explanation 
of  this  behaviour  will  be  evident  from  the  previous  section.  Both 
salts  are  soluble  in  sodium  thiosulphate  (p.  315).  They  are  acted  on 
by  light  in  the  same  way  as  the  chloride,  and  upon  this  depends  their 
use  in  photography. 

Silver  Fluoride,  AgF,  is  obtained  by  dissolving  silver  oxide  in 
hydrofluoric  acid  and  evaporating,  when,  according  to  the  conditions, 
the  compound  AgF,H2O,  colourless  quadratic  crystals,  or  AgF,2H2O, 
colourless  prisms,  is  obtained.  The  anhydrous  salt  is  amorphous 
and  generally  yellow,  but  the  pure  salt  should  presumably  be  colour- 
less. It  differs  from  the  other  silver  halides  in  being  readily  soluble 
in  water. 

Silver  Cyanide,  AgCN  —  This  salt  is  obtained  as  an  amorphous 
white  precipitate  when  excess  of  silver  nitrate  is  added  to  potassium 
cyanide.  It  is  readily  soluble  in  ammonia.  When  potassium  cyanide 
is  used  in  excess,  the  cyanide  dissolves  with  formation  of  a  complex 
compound,  KAg(CN)2.  This  compound  is  ionised  mainly  according 
to  the  equation 

KAg(CN)2^K-+Ag(CN)2', 

and  the  anion  is  very  slightly  ionised  thus  : 


so  that  the  Ag'  ion  concentration  is  very  minute. 

The  solution  of  potassium  silver  cyanide  is  used  in  electroplating. 
For  this  purpose  the  object  to  be  plated  is  immersed  in  a  solution  of 
the  cyanide,  and  forms  the  cathode,  the  anode  being  a  silver  plate. 
During  electrolysis  silver  is  dissolved  from  the  anode  and  deposits  on 
the  object  as  a  uniform  adherent  coating.  The  advantage  of  using 
the  cyanide  for  this  purpose  appears  to  depend  upon  the  small  Ag'  ion 


442     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

concentration,  as  highly  ionised  silver  salts,  e.g.  silver  nitrate,  give  a 
much  inferior  coating. 

Use  of  Silver  Halides  in  Photography—Modern  photo- 
graphy is  based  on  the  changes  undergone  by  silver  halides  on 
exposure  to  light.  The  "dry"  sensitive  plates  now  almost  exclusively 
used  are  glass  plates  coated  with  a  thin  film  of  gelatine  containing 
silver  bromide  in  suspension.  When  a  plate  of  this  kind  is  exposed 
in  a  camera,  the  silver  bromide  is  affected  in  some  parts  more  strongly 
than  in  others,  according  tP  the  relative  intensities  of  the  light  re- 
flected from  the  different  parts  of  the  object.  At  this  stage  no  visible 
change  has  taken  place  on  the  plate,  but  when  it  is  acted  on  by  a 
"  developer  " — for  example,  the  solution  of  a  reducing  agent  such  as 
ferrous  oxalate — silver  is  set  free  in  the  metallic  form  in  those  parts 
affected  by  the  light,  whilst  the  non-illuminated  silver  halide  is  un- 
affected. The  unaltered  halide  is  then  removed  by  treatment  with 
sodium  thiosulphate,  which  does  not  affect  the  deposited  silver — this 
process  is  known  as  fixing — and  the  result  is  the  so-called  "  negative  " 
on  which  the  illuminated  parts  are  dark  (owing  to  the  free  silver),  and 
the  unilluminated  parts  clear  (owing  to  the  removal  of  the  halide). 
Up  to  this  stage,  the  process  must  be  carried  through  in  the  absence 
of  daylight. 

The  negative  is  then  laid  on  a  paper  coated  with  a  sensitive  film 
(e.g.  silver  chloride  in  gelatine)  and  exposed  to  direct  sunlight.  Those 
parts  of  the  negative  where  the  silver  was  deposited  allow  very  little 
or  no  light  to  pass  through,  according  to  the  amount  of  silver  ;  those 
parts  free  from  silver  allow  all  the  light  to  pass.  After  "  fixing,"  a 
positive  which  reproduces  the  illumination  of  the  original  object  is 
obtained.  Of  the  silver  halides  the  bromide,  being  the  most  sensitive 
to  light,  is  most  largely  used  in  photography.  The  chemical  changes 
in  the  process  are  not  thoroughly  understood.  The  first  effect  of 
light  is  probably  to  form  a  "subhalide"  or  "photohalide,"  e.g.  Ag2Br. 

Silver  Sulphate,  Ag2SO4,  is  prepared  by  the  action  of  sulphuric 
acid  on  silver,  silver  oxide,  or  carbonate.  It  occurs  in  rhombic 
crystals,  isomorphous  with  sodium  sulphate. 

Silver  Nitrate,  AgNO3 — This  salt  is  obtained  by  the  action  of 
nitric  acid  on  silver.  It  occurs  in  colourless,  rhombic  crystals  which 
are  not  hygroscopic  ;  it  melts  at  208°.  It  is  extremely  soluble  in 
water;  at  o°  100  parts  of  water  dissolve  115  parts,  and  at  20°  215 
parts  of  the  salt.  When  heated  above  its  melting-point  it  decomposes 
into  silver  oxide,  oxygen  and  oxides  of  nitrogen  ;  above  300°  complete 
decomposition  into  metallic  silver  occurs. 


GOLD  443 

In  contact  with  organic  matter  and  other  reducing  agents  it  is 
reduced  (most  rapidly  on  exposure  to  light)  to  metallic  silver,  which 
is  usually  deposited  in  a  black,  finely-divided  form.  For  this  reason 
it  is  used  in  preparing  marking  inks  ;  it  is  also  employed  in  medicine 
as  a  caustic. 

From  solutions  of  silver  nitrate  in  ammonia  the  compound 
AgNO3,2NH3  separates  in  rhombic  crystals. 

Silver  Carbonate,  Ag2CO3,  is  obtained  as  a  light-yellow  powder 
by  double  decomposition  between  silver  nitrate  and  potassium  carbo- 
nate in  solution.  It  is  almost  insoluble  in  water. 

Tests  for  Silver — The  formation  of  a  white,  curdy  precipitate 
of  silver  chloride,  soluble  in  ammonia  but  insoluble  in  nitric  acid, 
when  hydrochloric  acid  or  a  chloride  is  added  to  the  solution  of  a 
silver  salt,  is  characteristic.  The  precipitation  of  dark  red  silver 
chromate,  Ag2CrO4,  when  potassium  chromate  is  added  to  the  solution 
of  a  silver  salt,  is  also  a  distinctive  test. 

GOLD 

Symbol,  Au.    Atomic  weight,  197.2. 

General  Characters — Gold  forms  two  series  of  salts,  the 
aurous  compounds,  of  the  type  AuX,  in  which  it  is  univalent,  and  the 
auric  compounds,  AuX3,  in  which  it  is  trivalent.  Auric  oxide,  Au2O3, 
has,  however,  acidic  as  well  as  basic  properties.  All  gold  compounds 
are  readily  reduced  to  metallic  gold  by  heat  or  by  reducing  agents. 

Occurrence — Gold  occurs  in  nature  mainly  in  the  free  con- 
dition in  veins  of  quartz  and  in  alluvial  deposits.  It  occurs  free  in 
small  amount  in  many  sulphide  ores,  e.g.  iron,  copper,  and  arsenical 
pyrites,  and  lead  and  zinc  sulphides.  In  the  combined  condition  it  is 
always  associated  with  tellurium,  and  also  usually  with  silver.  The 
more  important  ores  are  sylvanite,  (AujAg)Te,1  calaverite,  (Au,Ag)Te2, 
with  very  little  silver,  and  petzite,  (Au,Ag)Te2,  with  a  considerable 
proportion  of  silver.  The  chief  gold-producing  countries,  in  order  of 
relative  importance,  are  the  Transvaal,  the  United  States,  and 
Australia. 

Metallurgy  of  Gold — In  the  case  of  alluvial  deposits  gold  is 
separated  by  washing,  advantage  being  taken  of  its  high  specific 
gravity.  When  it  occurs  in  quartz  reefs,  as  in  the  Transvaal,  the 
rock  is  crushed,  reduced  to  a  fine  powder  by  stamping,  and  caused  to 

1  The  formula  (Au.Ag)Te  denotes  an  isomorphous  mixture  of  AuTe  and  AgTe 
in  varying  proportions. 


444     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

flow,  by  means  of  a  stream  of  water,  over  copper  plates  amalgamated 
with  mercury,  the  latter  retaining  the  gold.  The  amalgam  is  then 
scraped  off  and  the  mercury  removed  by  distillation. 

Besides  the  above  process,  two  methods  are  in  use  for  the  extraction 
of  gold  by  dissolution,  for  which  potassium  cyanide  and  chlorine 
respectively  are  used.  The  cyanide  method  is  carried  out  as  follows  : 
The  material  is  first  extracted  with  dilute  alkali  to  remove  certain 
impurities,  and  then  with  0.35  per  cent,  potassium  cyanide  solution 
with  free  access  of  air.  The  gold  is  then  precipitated  by  means  of  zinc 
or  by  electrolysis  ;  in  the  former  case  the  zinc  is  removed  by  roasting. 

In  order  that  gold  may  be  dissolved  by  potassium  cyanide,  the 
presence  of  oxygen  is  necessary.  The  equation  is  as  follows  :  — 


In  the  electrolytic  method  of  separation,  iron  anodes  and  lead 
cathodes  are  used.  The  latter  are  finally  subjected  to  cupellation  in 
order  to  obtain  pure  gold. 

The  chlorine  method  is  more  troublesome,  inasmuch  as  the  ore  has 
first  to  be  roasted  and  then  fused  with  common  salt  in  order  to 
convert  all  the  other  metals,  except  the  gold,  into  chlorides.  On 
treatment  with  chlorine,  the  gold  passes  into  solution  as  auric 
chloride,  AuCl3,  and  is  precipitated  as  sulphide  by  hydrogen  sulphide, 
or  as  metal  by  ferrous  sulphate  : 

2  AuCl3  +  6FeSO4-»2Au  +  2FeCl3+  2Fe2(SO4)3. 

Properties  —  Gold  is  a  bright  yellow,  rather  soft,  lustrous  metal, 
which  melts  at  1063°  :  density,  19.3  to  19.5  at  18°.  It  is  extremely 
malleable  and  ductile,  and  can  be  beaten  out  into  very  thin  sheets, 
which  are  green  by  transmitted  light.  It  is  a  very  good  conductor  of 
heat  and  electricity. 

Gold  is  not  acted  on  by  dry  or  moist  air  or  oxygen,  even  on 
heating,  and  is  not  affected  by  any  single  acid  (except  selenic  acid, 
p.  321).  Chlorine  readily  dissolves  it  with  formation  of  auric  chloride, 
AuCl3,  as  does  a  mixture  of  nitric  and  hydrochloric  acids  (aqua  regia), 
the  action  in  this  case  depending  upon  the  presence  of  free  chlorine 
(p.  235).  It  is  readily  dissolved  by  potassium  cyanide  solution  with 
access  of  air  (see  above).  Fused  alkalis  also  act  on  it. 

Compounds  of  gold  are  very  readily  reduced  to  the  metallic 
condition;  the  precipitated  gold  varies  in  appearance,  according  to 
the  nature  of  the  reducing  agent  and  the  conditions.  With  ferrous 
sulphate  or  arsenious  acid  a  brown  precipitate  is  obtained.  With 


GOLD  445 

other  reducing  agents,  such  as  formaldehyde  and  hydrazine,  gold  is 
obtained  in  colloidal  solution  showing  brilliant  colours,  such  as  purple, 
red,  and  blue. 

Alloys  —  Like  silver,  gold  is  too  soft  to  use  alone  for  coins  and 
other  articles,  and  is  always  alloyed  with  copper.  The  purity  or 
fineness  of  gold  is  expressed  in  carats,  pure  gold  being  24  carats. 
Gold  jewellery,  medals,  etc.,  are  14  to  18  carats,  that  is,  24  parts  of 
the  alloy  contains  14  to  18  parts  of  pure  gold.  The  British  gold 
coinage  contains  I  of  copper  to  1  1  of  gold  (22  carats),  that  of  the 
United  States  i  of  copper  to  9  of  gold. 

Aurous  Compounds  —  These  compounds,  which  are  of  the  same  type  as  the 
silver  salts,  are  characterized  by  the  readiness  with  which,  in  presence  of  water, 
they  yield  auric  compounds  and  the  metal  : 

3AuX->AuX3  +  2Au. 

Aurous  Oxide,  Au2O,  is  obtained  by  adding  an  alkali  hydroxide  to  the  solution 
of  an  aurous  salt.  It  is  a  dark  violet  powder,  which  decomposes  rapidly  into  gold 
and  oxygen  when  heated  to  250°. 

Aurous  Chloride,  AuCl,  is  obtained  as  a  yellowish-white  powder  by  heating  auric 
chloride  to  180°  : 


It  is  decomposed  by  water,  with  formation  of  auric  chloride  and  metallic  gold. 
With  hydrochloric  acid  and  the  alkali  chlorides  it  forms  complex  compounds  of 
the  type  H  AuCl2  and  KAuCla-  When  heated  to  230°,  it  is  completely  decomposed 
into  gold  and  chlorine. 

Aurous  Cyanide,  AuCN,  like  cuprous  cyanide  (cf.  p.  428),  is  obtained  under 
conditions  such  that  the  formation  of  auric  cyanide  might  be  expected.  It  forms 
yellow  microscopic  plates,  and  dissolves  in  excess  of  potassium  cyanide  to  form 
a  complex  salt,  KAu(CN)2,  potassium  aurocyanide,  which  is  used  in  gold-plating 
(cf.  silver-plating,  p.  441). 

Auric  salts  dissolve  in  potassium  cyanide  to  form  potassium  auricyanide, 
KAu(CN)4  (=  KCN,Au(CN)3),  also  used  in  gold-plating.  Auric  cyanide, 
Au(CN)3,  is  known  ;  it  occurs  in  large  colourless  plates  with  3H2O. 

Auric  Oxide,  Au2O3,  is  obtained  by  heating  auric  hydroxide,  Au(OH)3,  to 
140-150°.  At  155-165°  it  loses  oxygen,  and  finally  the  lower  oxide  remains. 
The  hydroxide,  Au(OH)3  (or  perhaps  AuO(OH)),  is  obtained  in  an  impure  state 
by  boiling  auric  chloride  with  potassium  hydroxide  ;  in  pure  condition  by  the 
action  of  magnesium  carbonate  on  auric  chloride,  the  magnesia  being  then 
removed  by  dilute  nitric  acid.  It  is  a  brown  powder,  soluble  in  excess  of 
potassium  hydroxide,  and  on  evaporating  the  solution  the  compound  KAuO2,3H2O, 
potassium  aurate,  separates  in  yellow  needles  : 

Au(OH)3+  KOH-»KAu02+2H2O. 
This  behaviour  shows  that  the  hydroxide  has  acidic  as  well  as  basic  properties. 


446     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Auric  Chloride,  AuCl3,  is  formed  when  gold  is  dissolved  in 
aqua  regia  or  in  chlorine  water,  and  is  obtained  as  a  dark  red  crystal- 
line mass  on  evaporating  the  solution  and  drying  carefully  at  1  50°. 
When  heated  to  180°  it  decomposes  into  aurous  chloride  and  chlorine  ; 
in  a  closed  space  an  equilibrium  between  the  three  substances  is 
established  : 


When  hydrochloric  acid  is  added  to  a  neutral  solution  of  gold 
chloride  the  colour  changes  to  yellow,  and  on  evaporating  the  com- 
pound HAuCl4  (  =  AuCl3,HCl),  chlorauric  acid,  separates  in  long, 
light-yellow  hygroscopic  needles.  Many  salts  are  derived  from 
chlorauric  acid,  such  as  KAuCl4,2H2O  and  NaAuCl4,2H2O.  In  solu- 
tions of  these  salts,  and  in  that  of  the  acid  itself,  the  gold  forms  part 
of  a  complex  anion,  AuCl4'. 

In  aqueous  solution  auric  chloride  is  partially  combined  to  form 
the  compound  HAuCl3(OH),  which  ionises  as  follows  :  — 


and  is  therefore  a  weak  acid.  It  is  clear  that  there  is  considerable 
analogy  between  this  compound  and  chlorauric  acid  ;  the  former  is 
derived  from  the  latter  by  the  substitution  of  OH  for  Cl. 

Gold  Sulphides  —  Auric  sulphide,  Au2S3,  cannot  be  obtained 
by  interaction  in  aqueous  solution,  as  it  is  immediately  decomposed 
by  water.  It  can,  however,  be  obtained  as  a  brown  powder  by  the 
action  of  hydrogen  sulphide  on  dry  lithium  chloraurate  at  -  10°,  the 
lithium  chloride  being  removed  from  the  product  by  treatment  with 
alcohol. 

When  hydrogen  sulphide  is  passed  into  a  boiling  solution  of  auric 
chloride,  aurous  sulphide,  Au2S,  mixed  with  sulphur,  is  obtained  as 
a  black  precipitate.  When  the  same  reaction  takes  place  at  room 
temperature  a  higher  sulphide,  probably  Au2S2,  is  obtained.  The 
sulphides  are  soluble  in  solutions  of  alkali  sulphides,  forming  thio- 
aurites,  e.g.  K3AuS2  or  3K2S,Au2S,  and  thioauraies,  KAuS2  or 
K2S,Au2S3.  (Cf.  thioarsenites  and  thioarsenates,  p.  519.) 

Purple  of  Cassius  is  a  substance  obtained  under  certain 
conditions  by  the  action  of  stannous  chloride  on  a  solution  of  auric 
chloride.  It  forms  a  brownish-purple  powder,  and  appears  to  be 
a  mixture  of  finely-divided  gold  with  stannic  acid  hydrogel  (p.  371), 
It  is  used  for  colouring  glass,  enamels,  etc. 

Tests  for  Gold  —  All  gold  compounds  finally  yield  the  metal 
on  heating.  Gold  is  also  obtained  as  a  brownish  precipitate  by  the 


GOLD  447 

action  of  reducing  agents  on  its  salts.  The  formation  of  purple  of 
Cassitis  is  also  characteristic. 

General  Characters  of  the  Sub-Group  and  Summary 

— The  resemblance  between  the  members  of  this  group  among 
themselves  is  considerably  less  than  in  other  groups  so  far 
considered.  The  metals  melt  at  high  temperatures,  they  are  not 
affected  by  water,  they  have  very  little  affinity  for  oxygen,  and  are 
among  the  least  electro-positive  of  the  metals.  The  small  affinity  is 
illustrated  by  the  fact  that  the  salts  are  readily  reduced  to  the  metal 
by  heating  or  by  the  action  of  reducing  agents.  They  resemble  each 
other  fairly  closely  in  their  univalent  -  ous  compounds  ;  the  univalent 
halides  are  white,  very  slightly  soluble  in  water,  form  complex  com- 
pounds with  hydrochloric  acid  of  the  type  HMC12,  and  the  cuprous 
and  silver  halides  are  soluble  in  ammonia  with  formation  of  com- 
pounds of  the  type  M(NH3)2C1.  A  striking  difference  in  behaviour  is 
that  silver  forms  only  univalent  compounds,  copper  forms  stable 
bivalent  compounds,  and  gold  tervalent  compounds. 

Contrary  to  the  behaviour  of  the  alkali  sub-group  (p.  422)  the 
electropositive  character  of  the  copper  sub-group  diminishes  with 
increasing  atomic  weight  (p.  434),  but  it  is  difficult  to  compare  the 
basic  character  within  the  latter  sub-group  on  account  of  the  existence 
of  different  series  of  salts.  Silver  hydroxide  is  a  much  stronger  base 
than  cupric  hydroxide  or  auric  hydroxide. 

The  contrast  in  behaviour  between  the  members  of  the  copper  and 
of  the  alkali  sub-group  is  very  marked.  Almost  the  only  resemblance 
is  that  all  function  as  univalent  elements,  the  crystals  of  cuprous, 
silver  and  sodium  chloride  belong  to  the  regular  system,  and  some 
corresponding  salts  of  silver  and  sodium  are  isomorphous.  The 
differences  can  readily  be  summarized  from  the  above,  and  the 
summary  of  the  behaviour  of  the  alkalis  (p.  422)  already  given.  The 
difference  in  electro-positive  character,  the  alkalis  being  very  strong 
and  the  members  of  the  copper  group  (except  silver)  very  weak  bases, 
is  important.  Further,  the  alkalis  have  no  tendency,  like  the  members 
of  the  copper  group,  to  form  complex  ions.  We  have  seen  that  the 
latter  metals  may  enter  both  into  complex  cations,  e.g.  Cn(NH3)2') 
Cu(NH3)4",  Ag(NH3)2-,  as  well  as  into  complex  anions,  e.g.  Ag(CN)2', 
Cu(CN)3",  AuCl/,  Au(CN)/. 

In  the  readiness  with  which  they  are  reduced  to  the  metallic 
condition  and  their  small  affinity  for  oxygen,  the  members  of  this 
group  resemble  the  platinum  group.  Copper  in  the  divalent  con- 
dition shows  considerable  resemblance  to  zinc,  divalent  nickel  and 
divalent  iron. 


CHAPTER  XXVIII 

ELEMENTS   OF   GROUP    II.,  SUB-GROUP   A 

• 

THIS  sub-group  comprises  the  following  three  metals,  which  are 
known  as  the  metals  of  the  alkaline  earths  : 

Atomic  Weight. 

Calcium  (Ca) 40.09 

Strontium  (Sr) «  87.63 

Barium  (Ba) 137-37 

The  members  of  this  family  are  invariably  divalent  in  their  compounds. 
The  hydroxides  are  strong  bases,  almost  as  strong  as  the  alkali 
hydroxides,  so  that  the  halides  are  stable  towards  water,  and  there  is 
little  or  no  tendency  to  the  formation  of  basic  salts.  The  metals 
themselves  are  soft,  readily  oxidize  in  the  air,  'and  are  acted  on  by 
water  at  room  temperature. 

CALCIUM 

Symbol,  Ca.     Atomic  weight  =40.09. 

Occurrence — In  the  combined  state  calcium  is  very  widely  distri- 
buted in  nature.  As  the  carbonate  it  occurs  in  enormous  amount  in  the 
form  of  chalk,  marble,  limestone,  and  coral.  As  sulphate  it  occurs  in 
gypsum  and  selenite^  CaSO4,2H2O,  and  as  anhydrite,  CaSO4.  As  fluo- 
ride, CaF2,  it  constitutesy?#0r.y/<2r/  and  occurs  as  phosphate,  Ca3(PO4)2, 
in  phosphorite.  Calcium  silicate,  CaSiO3,  is  a  constituent  of  many 
rocks.  Calcium  salts  are  found  in  all  soils,  from  which  they  are  taken 
up  by  plants ;  also  in  most  natural  waters.  Bones  are  mainly  com- 
posed of  calcium  phosphate. 

Preparation  of  Metal — Calcium  is  obtained  in  a  nearly  pure 
state  by  heating  calcium  iodide  with  metallic  sodium,  the  excess 
of  sodium  being  finally  removed  by  treatment  with  anhydrous  alcohol 
(Moissan). 

A  more  convenient  method,  now  used  exclusively  for  commercial 
purposes,  depends  upon  the  electrolysis  of  the  fused  chloride.  The 
salt  is  fused  in  a  vessel  of  graphite,  the  walls  of  which  form  the  anode  ; 

448 


CALCIUM  449 

and  an  iron  rod,  just  touching  the  surface  of  the  electrolyte,  is  used 
as  cathode.  As  electrolysis  proceeds  the  light  metal  collects  at 
the  end  of  the  cathode,  and  the  latter  is  slowly  raised  in  ordef 
that  the  metal  may  solidify  out  of  contact  with  the  fused  electro- 
lyte. In  this  way,  by  progressively  raising  the  cathode,  a  long 
irregular  rod  of  metal  may  be  built  up  without  interruption  of  the 
electrolysis. 

Properties — Calcium  is  a  silvery  white  metal  somewhat  harder 
than  lead,  its  density  is  1.52,  and  it  melts  at  800°.  It  turns  yellow  on 
the  surface  on  exposure  to  air,  probably  owing  to  the  formation  of 
nitride,  Ca3N2.  It  burns  vigorously  when  heated  in  air  or  oxygen, 
oxide  and  nitride  being  formed.  It  decomposes  water  slowly  at  room 
temperature,  hydrogen  being  liberated. 

Calcium  Hydride,  CaH2,  is  obtained  by  passing  hydrogen 
over  calcium  at  a  red  heat.  It  is  a  white  powder,  which  is  vigorously 
acted  on  by  water  with  liberation  of  hydrogen. 

Calcium  Oxides — Two  oxides  of  calcium  are  known  :  the 
normal  oxide,  CaO,  and  the  dioxide,  CaO2. 

Calcium  Oxide  (lime,  quicklime),  CaO,  is  prepared  by  heat- 
ing calcium  carbonate  to  a  temperature  exceeding  800°  under  such 
conditions  that  the  carbon  dioxide  set  free  is  continually  removed  : 

CaCO3->CaO  +  CO2. 

On  the  commercial  scale  this  process  is  carried  out  by  heating  lime- 
stone in  brick  kilns. 

The  calcium  oxide  obtained  by  burning  limestone  is  a  white  amor- 
phous powder  which  is  infusible  in  the  oxyhydrogen  flame  (lime- 
light), but  readily  fuses  in  the  electric  furnace.  It  unites  with  water 
with  the  evolution  of  much  heat,  calcium  hydroxide,  Ca(OH)2,  being 
formed  ;  in  this  process  lumps  of  quicklime  crumble  to  powder.  This 
is  known  as  the  slaking  of  lime,  and  the  product  is  termed  slaked 
lime.  On  account  of  its  great  affinity  for  water,  lime  is  often  used  for 
drying  gases  (p.  215)  and  for  removing  the  last  traces  of  water  from 
liquids. 

Calcium  Hydroxide,  Ca(OH)2,  is  a  white,  amorphous,  hygro- 
scopic powder.  When  it  dissolves  in  water  heat  is  given  out,  and 
therefore  the  solubility  diminishes  with  rise  of  temperature.  At  20° 
TOO  grams  of  water  dissolve  0.126  grams,  at  50°  0.098  grams,  at  100° 
0.060  grams  of  the  base.  The  solution  is  termed  lime-water.  Water 
containing  a  large  excess  of  calcium  hydroxide  in  suspension  is  termed 
milk  of  lime.  Milk  of  lime,  lime-water,  and  calcium  oxide  itself  readily 
29 


450     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

absorb  carbon  dioxide  from  the  air,  with  formation  of  calcium 
carbonate. 

Calcium  Peroxide,  CaO2,  separates  as  the  octahydrate, 
CaO2,8H2O,  when  hydrogen  or  sodium  peroxide  is  added  to  lime-water. 
On  heating  to  130°  it  becomes  anhydrous,  and  at  a  red  heat  decom- 
poses into  calcium  oxide  and  oxygen. 

Mortar  and  Cement  —  Mortar  consists  of  lime  and  sand  made 
into  a  paste  with  water.  The  setting  is  due  to  the  escape  of  water 
and  the  absorption  of  carbon  dioxide  to  form  calcium  carbonate, 
which  sets  into  a  solid  mass  with  the  sand.  The  latter  renders  the 
mass  porous,  thus  facilitating  the  entry  of  the  carbon  dioxide.  No 
calcium  silicate  is  formed  till  after  the  lapse  of  many  years,  so  that 
this  substance  plays  no  part  in  the  process  of  hardening. 

Ordinary  Cement  is  a  mixture  of  lime  (50  to  60  per  cent.),  silica 
(25  per  cent.),  and  aluminium  oxide  (8  to  10  per  cent.).  It  is  made 
by  burning  a  mixture  of  limestone  and  clay.  In  some  localities  mix- 
tures of  limestone  and  aluminium  silicates  occur  which  yield  cement 
directly  on  burning.  Cement  has  the  advantage  over  mortar  of  set- 
ting to  a  hard  mass  even  under  water  (hydraulic  cement).  Portland 
cement  is  of  similar  composition  to  ordinary  cement.  It  is  prepared 
by  burning  at  a  high  temperature  a  mixture  in  definite  proportions  of 
limestone  and  clay  rich  in  silica  ;  the  hard  mass  or  clinker  is  then 
reduced  to  a  fine  powder.  The  hardening  of  cement  is  by  no  means 
understood.  According  to  Le  Chatelier  it  depends  mainly  on  the 
change  of  a  basic  calcium  silicate  by  absorption  of  water  to  the  normal 
hydrated  silicate  and  calcium  hydroxide  : 


Calcium  Carbonate,  CaCO3,  is  perhaps  the  most  familiar  ex- 
ample of  a  dimorphous  substance,  being  met  with  in  the  two  crystal- 
line forms  :  calcite,  which  generally  occurs  in  rhombohedral  crystals 
belonging  to  the  hexagonal  system  ;  and  aragonite,  in  crystals  belong- 
ing to  the  orthorhombic  system.  Chalk,  marble,  limestone,  Iceland 
spar,  etc.,  belong  to  the  calcite  modification  ;  aragonite  is  of  somewhat 
rare  occurrence  in  nature.  When  precipitated  from  solution  by  inter- 
action of  a  soluble  calcium  salt  and  sodium  carbonate  at  room 
temperature,  calcium  carbonate  is  amorphous  ;  but  on  standing  in 
contact  with  the  mother  liquor  changes  to  minute  crystals  of  calcite. 
When,  on  the  other  hand,  precipitation  occurs  from  hot  solution, 
aragonite  is  obtained.  Calcite  appears  to  be  the  stable  form  above 


CALCIUM  451 

o°  under  all  conditions  ;  the  apparent  stability  of  aragonite  is  due  to  its 
slow  rate  of  change  (p.  174). 

Calcium  carbonate  is  practically  insoluble  in  water ;  but  is  soluble 
in  water  containing  carbon  dioxide.  In  this  case  soluble  calcium 
bicarbonate,  CaH2(CO3)2,  is  doubtless  present  in  the  solution  : 

CaC03  +  H2O  +  CO2^CaH2(CO3)2. 

The  reaction  is  reversible ;  on  boiling  carbon  dioxide  escapes  and 
calcium  carbonate  is  reprecipitated.  The  deposit  or  "fur"  in  kettles 
and  steam  boilers  is  calcium  carbonate  which  has  fallen  out  of  solution 
owing  to  the  removal  of  the  carbon  dioxide  on  boiling. 

Hardness  of  Water — As  already  indicated  (p.  341),  the  hard- 
ness of  water  is  due  almost  entirely  to  the  presence  of  calcium  and 
magnesium  salts  in  solution,  chiefly  as  carbonates  and  sulphates.  A 
distinction  is  drawn  between  temporary  hardness,  due  to  the  carbon- 
ates held  in  solution  by  carbon  dioxide  and  which  is  removed  by 
boiling,  and  permanent  hardness,  due  to  sulphates  and  chlorides, 
which  cannot  be  removed  by  boiling. 

The  hardness  of  water  is  chiefly  noticeable  in  its  action  on  soap. 
A  soft  water,  containing  little  dissolved  salts,  readily  forms  a  lather 
with  soap  ;  but  hard  waters  use  up  a  large  quantity  of  soap  before  a 
lather  is  obtained.  The  hardness  of  water  (including  both  temporary 
and  permanent  hardness)  can,  in  fact,  be  estimated  by  adding  to  water 
a  standard  soap  solution  from  a  burette  till  the  point  is  reached  at 
which  a  lather  is  formed  on  shaking.  Temporary  hardness  can  be 
removed  by  boiling,  as  already  indicated,  and  also  by  adding  to  the 
•  solution  sufficient  calcium  oxide  to  combine  with  the  carbon  dioxide 
present,  when  all  the  carbonate  is  precipitated  : 

Ca(HCO8)2  +  CaO-»2CaCO8  +  H2O. 

Both  temporary  and  permanent  hardness  can  be  removed  by  the 
addition  of  sodium  carbonate  : 

CaSO4+Na2CO3-->CaCO3J+Na2SO4. 

Calcium  Chloride,  CaCl2,  occurs  as  tachydrite  (2MgCl2-CaCl2. 
I2H2O)  in  the  Stassfurt  deposits.  It  is  a  bye-product  in  certain 
technical  operations,  such  as  the  preparation  of  ammonia  from  its 
salts  (p.  214).  It  separates  from  solution  at  room  temperature  as 
the  hexahydrate,  CaCl2,6H2O,  in  hexagonal  crystals  ;  other  hydrates 
are  also  known.  The  hexahydrate  melts  at  30.2°.  Above  260°  all 
the  hydrates  change  to  the  anhydrous  salt ;  the  latter  melts  at  802°. 
Anhydrous  calcium  chloride  has  a  great  affinity  for  water,  and  is  used 
for  drying  gases  and  organic  liquids.  It  forms  a  compound  with 


452     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

ammonia,  CaCl2,8NH3,  and  cannot  therefore  be  used  for  drying 
this  gas. 

The  hexahydrate  and  ice  form  a  very  useful  freezing-mixture 
(p.  199),  the  temperature  (cryohydric  temperature)  falling  to  —55°. 

Chlorinated  Lime  (bleaching  powder) — This  substance 
is  prepared  by  the  prolonged  action  of  chlorine  on  damp  slaked  lime 
at  room  temperature.  The  reaction  may  be  represented  by  the  follow- 
ing equation  : 

Ca(OH)2+Cl2->Cl.Ca.OCl+H20, 

but  a  certain  amount  of  unaltered  calcium  hydroxide  is  always  present. 
It  was  formerly  supposed  that  bleaching  powder  was  a  mixture  of 
equivalent  amounts  of  calcium  chloride  and  calcium  hypochlorite, 
CaCl2,Ca(OCl)2,  but  conclusive  evidence  has  been  obtained  that  the 
solid  compound  contains  no  free  calcium  chloride.1  The  aqueous 
solution,  however,  appears  to  contain  a  mixture  of  chloride  and  hypo- 
chlorite : 

2Ca(OCl)Cl->CaCl2+Ca(OCl)2. 

Dilute  acids,  even  moist  carbon  dioxide,  liberate  practically  all  the 
chlorine  from  bleaching  powder  even  in  the  cold.  In  the  case  of 
hydrochloric  acid,  we  may  assume  that  hypochlorous  acid  is  first 
formed,  and  that  it  immediately  reacts  with  hydrochloric  acid  to  form 
water  and  chlorine  : 

Ca(OCl)2+2HCl->CaCl2  +  2HClO. 
HC1O  +  HC1->H2O  +  C12. 

Other  acids  presumably  liberate  both  hydrochloric  and  hypochlorous  % 
acids,  which  then  react  as  above. 

The  use  of  this  substance  in  bleaching  depends  upon  the  reactions 
just  considered.  The  material  to  be  bleached  is  dipped  first  in  a 
solution  of  bleaching  powder  and  then  into  a  dilute  acid  ;  the  effect 
is  due  mainly  to  free  hypochlorous  acid  (p*  91). 

The  "available"  chlorine  is  the  amount  set  free  from  bleaching 
powder  by  the  action  of  excess  of  a  weak  acid  such  as  acetic  acid.  It 
usually  amounts  to  30  to  36  per  cent.,  but  a  product  containing  over 
39  per  cent,  can  be  obtained. 

Calcium  Sulphate,  CaSO4,  occurs  naturally  in  the  anhydrous 
form  as  anhydrite  (rhombic  crystals),  and  as  hydrate  in  gypsum, 
CaSO4,2H2O  (monoclinic  crystals).  Alabaster  and  selenite  are  forms 

1  The  main  evidence  in  favour  of  this  view  is  that  bleaching  powder,  unlike 
calcium  chloride,  is  not  deliquescent ;  and  further,  alcohol,  though  a  good  solvent 
for  calcium  chloride,  does  not  extract  the  latter  from  bleaching  powder. 


CALCIUM  453 

of  gypsum.  When  calcium  sulphate  is  prepared  by  double  decom- 
position, the  dihydrate  is  obtained. 

When  gypsum  is  heated  in  an  open  vessel  at  98°,  it  loses  water  and 
the  so-called  half-hydrate,  (CaSO4)2,H2O,  is  obtained  ;  on  prolonged 
heating  at  107°  to  108°,  more  readily  at  higher  temperatures,  all  the 
water  is  driven  off,  and  one  form  of  the  anhydrous  salt,  the  so-called 
"soluble  anhydrite"  is  obtained.  When  either  the  half-hydrate  or 
soluble  anhydrite  is  treated  with  water,  rehydration  takes  place,  and 
the  mixture  sets  to  a  hard  mass. 

Plaster  of  Paris  consists  mainly  of  the  half-hydrate,  and  is 
made  by  heating  gypsum  in  kilns  at  a  temperature  very  little  over 
130°.  The  setting  is  represented  as  follows  : — 

(CaSO4)2,H2O  +  3H2O-»2(CaSO4,2H2O). 

When  gypsum  is  heated  for  some  time  above  200°,  a  second  modifi- 
cation of  anhydrite  is  formed,  which  sets  very  slowly  with  water. 
Gypsum  thus  treated  is  said  to  be  dead  burnt,  and  the  product  is  use- 
less for  making  casts.  Natural  anhydrite  does  not  set  with  water. 

Calcium  sulphate  is  only  slightly  soluble  in  water.  The^olubility 
increases  slowly  with  rise  of  temperature  up  to  35°,  and  then  diminishes 
(cf.  p.  84).  The  amount  dissolved  depends  somewhat  on  the  size  of 
the  particles  ;  the  finer  the  state  of  division  the  greater  is  the  solu- 
bility. This  rule  is  a  quite  general  one,  but  the  difference  is  only 
appreciable  for  substances  of  slight  solubility.  At  o°  100  grams  of 
water  take  up  about  0.195  grams,  at  38°  0.221  grams,  and  at  99° 
o.i  80  grams  of  the  salt. 

Calcium  Phosphates — Tricalcium  orthophosphate,  Ca3(PO4)2, 
is  the  chief  inorganic  constituent  of  bones.  It  also  occurs  naturally 
in  immense  quantities  as  phosphorite  and  osteolite,  and  along  with 
calcium  fluoride  in  apatite^  3Ca3(PO4)2,CaF2.  As  it  is  practically 
insoluble  in  water,  it  is  at  once  precipitated  when  solutions  contain- 
ing Ca"  and  PO4'"  ions  respectively  are  mixed,  e.g.  by  adding  am- 
monia and  sodium  phosphate  to  a  solution  of  calcium  chloride. 
Although  very  slightly  soluble  in  water,  it  is  soluble  in  acids,  even 
carbonic  acid,  and  in  solutions  of  salts,  more  particularly  ammonium 
salts  and  the  chlorides  and  nitrates  of  the  alkalis.  These  facts  are  of 
great  importance  in  facilitating  the  taking  up  of  calcium  phosphate  by 
plants  from  the  soil,  as  plants  can  only  take  up  soluble  substances. 

When  acted  on  by  the  requisite  quantity  of  sulphuric  acid,  mono- 
calcium  phosphate,  CaH4(PO4)2,  and  calcium  sulphate  are  formed  : 


454    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

This  mixture  is  known  as  superphosphate  of  lime,  and  is  largely  used 
as  a  manure.  Monocalcium  phosphate  has  the  advantage  of  being 
readily  soluble  in  water. 

The  intermediate  phosphate,  dicaldum  phosphate,  Ca2H2(PO4)2  or 
CaHPO4,  is  obtained  by  decomposing  the  monocalcium  salt  with 
water  : 

CaH4(PO4)2-»CaHPO4>J,  +  H3PO4. 

It  is  more  soluble  in  water  than  the  triphosphate,  but  much  less 
soluble  than  the  monocalcium  phosphate. 

Calcium  Carbide,  CaC2,  is  prepared  by  heating  lime  with 
carbon  in  the  electric  furnace  (p.  331)  : 

CaO  +  3C->CaC2  +  CO. 

Pure  calcium  carbide  is  colourless  ;  the  dark  colour  of  the  com- 
mercial article  is  due  to  impurities.  It  is  used  commercially  in  the 
preparation  of  acetylene. 

Calcium  Cyanamide,  CaNCN,  has  already  been  referred  to 
in  connexion  with  the  utilization  of  atmospheric  nitrogen  (p.  236). 

Calcium  Sulphide,  CaS,  is  obtained  by  reducing  calcium 
sulphate  with  carbon.  When  pure,  it  is  a  white,  amorphous  powder 
It  is  very  slightly  soluble  as  such  in  water,  but  undergoes  hydrolytic 
decomposition  with  formation  of  calcium  hydrosulphide,  Ca(SH)2, 
and  calcium  hydroxide,  the  former  of  which  is  readily  soluble  in 
water  : 

2O->Ca(SH)2  +  Ca(OH)2. 


For  this  reason  no  precipitate  is  obtained  when  ammonium  sulphide 
is  added  to  a  calcium  salt  in  solution  ;  owing  to  hydrolysis  only  the 
soluble  hydrosulphide  is  formed. 

Ordinary  calcium  sulphide,  as  well  as  the  sulphides  of  strontium 
and  barium,  after  exposure  to  light,  remain  luminous  in  the  dark. 
For  this  reason  they  are  used  in  making  luminous  paint.  The  pure 
substances  do  not  show  this  property,  which  is  therefore  connected 
with  the  presence  of  small  amounts  of  impurities. 

Calcium  Silicate.  Glass—  The  metasilicate^  CaSiO3,  occurs 
naturally  as  wollastonite,  and  can  be  prepared  by  fusing  together 
calcium  oxide  and  silica  in  the  electric  furnace.  It  occurs  in  crystals 
insoluble  in  water,  and  the  temperature  has  to  be  raised  to  1400°  in 
order  to  fuse  it. 

The  different  kinds  of  glass  are  mixtures  of  silicates  of  calcium  or 
lead  with  alkali  silicates,  and  contain  excess  of  silica.  Ordinary  soft 


STRONTIUM  ,       455 

glass  (window  gla:s)  is  made  by  fusing  together  sodium  carbonate, 
limestone,  and  sand.  It  is  readily  fusible.  Crown  or  Bohemian 
glass  contains  potassium  instead  of  sodium  silicate,  and  is  less 
readily  fusible  than  ordinary  glass.  Flint  glass  is  a  mixture  of  lead 
and  potassium  silicates.  It  is  readily  fusible,  and  has  a  high  refrac- 
tive index.  The  colours  of  certain  kinds  of  glass  are  due  to  traces 
of  impurities.  The  green  colour  of  the  common  glass  used  for 
making  bottles  and  window-panes  is  due  to  silicate  of  iron.  The 
appearance  of  milk-glass  is  due  to  the  presence  of  finely-divided 
calcium  phosphate. 

Glass  has  no  definite  crystalline  properties,  and  when  heated 
gradually  softens  without  showing  a  definite  melting-point ;  it  must 
therefore  be  regarded  as  an  amorphous  substance,  and  is  in  reality 
a  highly  supercooled  liquid  (p.  69).  When  kept  at  a  high  tempera- 
ture for  a  long  time,  some  of  the  silicates  may  separate  in  crystalline 
form  ;  the  glass  is  then  said  to  be  "  devitrified,"  and  is  very  brittle. 

As  the  alkali  silicates  are  soluble  in  water  (p.  370),  it  is  not  sur- 
prising that  water  (especially  hot  water)  dissolves  out  small  amounts 
of  alkali  from  glass.  This  may  be  shown  by  grinding  up  some  soft 
glass  with  water  in  a  mortar  and  adding  a  few  drops  of  phenol- 
phthalein,  when  the  solution  turns  pink. 

Tests  for  Calcium — Calcium  compounds  give  a  brick-red 
colour  to  the  Bunsen  flame.  The  most  characteristic  wet  test  is  the 
precipitation  of  calcium  oxalate  by  double  decomposition  ;  this  salt 
is  insoluble  in  acetic,  but  readily  soluble  in  hydrochloric  acid. 

STRONTIUM 

Symbol,  Sr.     Atomic  Weight =87. 63. 

The  compounds  of  strontium  are  very  similar  to  those  of  calcium, 
and  may  therefore  be  discussed  very  briefly. 

Occurrence — Strontium  is  much  less  widely  distributed  than  the 
other  members  of  the  group.  The  chief  naturally  occurring  com- 
pounds are  celestine,  SrSO4,  and  strontianite^  SrCO3. 

Preparation  of  Metal — Strontium,  like  calcium,  is  usually 
prepared  by  electrolysis  of  the  fused  chloride. 

Properties — Strontium  is  a  soft,  silvery-white  metal  of  density 
2.55  ;  it  melts  about  800°.  It  is  acted  on  by  water  and  by  dilute  acids 
at  room  temperature  more  vigorously  than  calcium,  and  when  heated 
in  oxygen  it  burns  vigorously  to  the  oxide.  When  heated  in  hydrogen, 
strontium  hydride,  SrH2,  is  formed. 


456    A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

Preparation  of  Strontium  Compounds — Strontium  com- 
pounds are  prepared  by  dissolving  the  carbonate,  strontianite,  in  the 
appropriate  acid,  or  by  reducing  the  sulphate  to  sulphide  by  means  of 
carbon,  and  treating  the  latter  compound  with  acids. 

Strontium  Oxide,  like  the  other  oxides  of  this  sub-group,  is 
obtained  by  strongly  heating  the  hydroxide,  carbonate,  or  nitrate.  It 
combines  readily  with  water  to  form  the  hydroxide,  Sr(OH)2;  the 
latter  separates  from  aqueous  solution  in  the  hydrated  form  as 
Sr(OH)2,8H2O  in  tetragonal  crystals.  Strontium  hydroxide  is  more 
soluble  in  water  than  calcium  hydroxide  (see  table,  p.  392),  and  the 
solubility  increases  regularly  with  the  temperature.  It  forms  a  com- 
pound with  sugar  which  is  insoluble  in  water,  but  is  readily  decom- 
posed by  carbon  dioxide,  and  for  this  reason  is  used  in  the  sugar 
industry. 

Strontium  Chloride,  SrCl2,  separates  from  solution  as 
SrCl2,6H2O  in  hexagonal  crystals,  isomorphous  with  calcium  chloride 
hexahydrate.  Unlike  barium  chloride,  it  is  somewhat  soluble  in 
alcohol. 

Strontium  Sulphate,  SrSO4,  obtained  by  double  decomposi- 
tion, is  less  soluble  in  water  than  calcium  sulphate,  but  more  soluble 
than  barium  sulphate.  At  20°  100  grams  of  water  take  up  0.1479 
grams,  at  80°  0.1688  grams  of  the  salt. 

Strontium  Nitrate,  Sr(NO3)2,  separates  from  warm  aqueous 
solutions  as  the  anhydrous  salt  in  octahedral  crystals  ;  from  cold 
solutions  as  Sr(NO3)2,4H2O  in  monoclinic  crystals.  It  is  practically 
insoluble  in  alcohol,  and  this  property  is  taken  advantage  of  in 
separating  strontium  from  calcium,  as  calcium  nitrate  is  soluble  in 
alcohol.  This  and  other  strontium  salts  are  used  in  pyrotechny,  owing 
to  the  deep  crimson  colour  they  impart  to  flame. 

Tests — Strontium  compounds  are  characterized  by  the  crimson 
flame,  very  similar  to  that  due  to  lithium.  The  spectra  of  the  two 
metals  are,  however,  quite  distinct ;  that  of  strontium  has  a  number 
of  bands  in  the  red  and  a  characteristic  blue  line. 

BARIUM 

Symbol,  Ba.     Atomic  Weight =137. 37. 

Occurrence — The  principal  natural  compounds  of  barium  are 
ivitherite,  BaCO3,  and  heavy  spar,  BaSO4. 

Preparation  of  Metal — The  preparation  of  pure  barium  has 
proved  a  matter  of  considerable  difficulty.  The  best  results  appear 
to  have  been  obtained  by  warming  a  concentrated  solution  of  barium 


BARIUM  457 

chloride  with  sodium  amalgam.  From  the  resulting  barium  amalgam 
the  mercury  is  removed  by  placing  it  in  a  tube  and  gradually 
increasing  the  temperature.  Finally,  at  1150°,  pure  barium  distils 
over. 

Properties  —  Barium  is  a  soft  and,  when  pure,  presumably  a 
silvery-white  metal,  which  melts  below  1000°  ;  its  density  is  3.75.  It 
becomes  oxidized  very  readily  in  the  air,  and  decomposes  water 
vigorously  at  room  temperature. 

Preparation  of  Compounds  —  Barium  compounds,  like  those 
of  strontium,  can  be  obtained  by  dissolving  the  carbonate  in  the 
appropriate  acid,  or  by  reducing  the  sulphate  to  the  sulphide  by 
heating  with  charcoal  and  then  dissolving  the  sulphide  in  acids.  A 
further  method,  starting  with  the  very  insoluble  sulphate,  is  to  fuse  it 
with  sodium  carbonate  : 


If  a  considerable  excess  of  the  carbonate  is  used,  the  action  is  almost 
complete  in  the  direction  of  the  upper  arrow.  The  sodium  sulphate 
and  excess  of  carbonate  are  removed  by  boiling  the  product  with 
water. 

Barium  Oxides  —  Two  oxides  of  barium  are  known,  the  monoxide, 
BaO,  and  the  peroxide,  BaO2. 

Barium  Oxide,  BaO,  is  usually  prepared  by  heating  the  nitrate. 
It  cannot  conveniently  be  obtained  by  heating  the  carbonate  alone, 
owing  to  the  very  high  temperature  at  which  the  latter  decomposes  ; 
but  when  the  carbonate  is  mixed  with  carbon,  the  reaction  takes  place 
at  a  much  lower  temperature  : 

BaCO3  -f  C->BaO  +  2CO. 

Barium  oxide  forms  a  white  amorphous  powder,  which  fuses  in  the 
electric  furnace  at  a  lower  temperature  than  calcium  oxide.  When 
heated  in  air  it  combines  with  oxygen  to  form  the  peroxide.  It  com- 
bines very  vigorously  with  water  to  form  the  hydroxide,  Ba(OH)2,  a 
large  amount  of  heat  being  given  out. 

Barium  Hydroxide,  Ba(OH)2,  obtained  as  described  above, 
separates  from  aqueous  solution,  at  all  temperatures  between  20°  and 
109°,  in  tetragonal  crystals  as  the  octahydrate,  Ba(OH)2,8H2O.  The 
octahydrate  effloresces  in  the  air,  with  formation  of  the  monohydrate. 
The  anhydrous  compound  melts  at  a  red  heat  without  giving  up 
water.  , 

Barium  hydroxide  is  much  more  soluble  in  water  than  the  other 


458     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

bases  of  this  group — the  solution  is  called  baryta-water.  At  10°  100 
grams  of  water  take  up  2.5  grams,  at  20°  4.3  grams,  and  at  40°  8.2 
grams  of  Ba(OH)2. 

Barium  Peroxide,  BaO2,  is  obtained  by  heating  barium  oxide 
in  oxygen,  or  air  free  from  carbon  dioxide,  at  temperatures  in  the 
neighbourhood  of  500°.  Raising  the  temperature  favours  the  reverse 
decomposition.  It  is  also  precipitated  in  crystalline  form  as 
BaO2,8H2O,  when  hydrogen  peroxide  is  added  to  baryta  water. 

Barium  peroxide  is  a  greyish  powder,  only  very  slightly  soluble  in 
water.  By  the  action  of  dilute  acids  hydrogen  peroxide  is  formed 
(p.  138). 

BaO2  +  H2SO4->BaSO4  +  H2O2. 

Barium  Chloride,  BaCl2,  prepared  by  the  general  methods 
(P-  39°X  separates  from  solution  in  colourless,  monoclinic  crystals 
as  BaCl2,2H2O.  Unlike  calcium  chloride  hexahydrate,  it  is  not 
deliquescent ;  it  is  readily  dehydrated  on  heating  to  100°.  It  is  very 
soluble  in  water,  and  is  precipitated  from  a  concentrated  solution  by 
adding  hydrochloric  acid  (cf.  p.  439). 

Barium  Sulphate,  BaSO4,  prepared  by  double  decomposition, 
is  a  white  powder,  which  can  be  obtained  in  the  amorphous  form  and 
also  as  minute  crystals.  It  is  practically  insoluble  in  water  (0.235 
milligrams  in  100  grams  of  water  at  18°)  and  in  dilute  acids.  It  is 
readily  soluble  (to  the  extent  of  10  to  12  per  cent.)  in  concentrated 
sulphuric  acid,  doubtless  owing  to  the  formation  of  the  acid  sulphate, 
Ba(HSO4)2,  but  is  reprecipitated  on  the  addition  of  water. 

Barium  Nitrate,  Ba(NO3)2,  obtained  by  the  general  methods, 
separates  from  aqueous  solution  at  room  temperature  in  the  anhydrous 
form  (octahedral  crystals).  It  is  readily  soluble  in  water.  It  is  used 
in  pyrotechny  for  the  production  of  green  flame. 

Tests  for  Barium — Barium  salts  impart  a  green  colour  to  the 
Bunsen  flame,  and  the  spectrum  is  characterized  by  the  presence  of 
certain  lines  in  the  green  and  orange.  In  the  wet  way,  the  formation  of 
the  sulphate,  which  is  always  obtained,  on  account  of  its  very  small 
solubility  product,  when  Ba"  and  SO4"  ions  are  brought  together, 
even  in  minute  amount,  is  the  most  important  test.  It  differs  from  the 
other  two  metals  of  the  sub-group  in  giving  a  yellow  precipitate  of 
barium  chromate  with  potassium  chromate,  insoluble  in  acetic  acid. 


METALS  OF  THE  ALKALINE  EARTHS 


459 


COMPARISON  OF  THE  ALKALINE  EARTH  METALS  AND 
SUMMARY 

As  is  evident  from  the  foregoing,  there  is  a  very  close  resemblance 
in  the  behaviour  of  the  corresponding  compounds  of  the  members  of 
this  group,  and,  further,  the  properties  vary  regularly  with  increase  of 
atomic  weight. 

The  metals  themselves  are  almost  as  electro-positive  as  the  alkali 
metals  ;  they  quickly  oxidize  in  the  air  and  decompose  water  at  room 
temperature.  They  impart  characteristic  colours  to  the  Bunsen  flame- 
They  are  all  divalent  elements.  Their  more  important  physical  pro- 
perties are  shown  in  the  table — 


Ca. 

Sr. 

Ba. 

Atomic  weight 

40.09 

87-63 

137-37 

Density 
Melting-point         .         . 

1.52 
800° 

2-55 
about  800° 

3-75 
about  800° 

Atomic  volume 

26.4 

33-7 

36.6 

The  chemical  properties  are,  of  course,  determined  by  their  strong 
electro-positive  character,  as  illustrated  by  the  fact  that  the  chlorides 
are  not  hydrolyzed,  and  that  there  is  very  little  tendency  to  form  basic 
salts.  Barium  appears  to  be  the  most  strongly  basic,  as  is  shown,  for 
example,  by  the  fact  that  the  hydroxide  can  be  fused  without  decom- 
position, whilst  calcium  hydroxide  readily  splits  up  into  water  and  the 
oxide  on  heating ;  the  relatively  great  stability  of  barium  carbonate 
also  supports  this  conclusion. 

The  alkaline  earths  differ  markedly  from  the  alkalis  with  regard  to 
the  small  solubility  of  the  hydroxides,  carbonates,  phosphates  and 
sulphates,  but  in  this  respect,  as  we  have  seen,  lithium  approximates 
somewhat  to  the  alkaline  earth  metals  (p.  422).  The  splubility  of  the 
hydroxides  and  carbonates  increases,  that  of  the  sulphates  decreases 
with  increasing  atomic  weight. 


T 


CHAPTER  XXIX 
ELEMENTS  OF  GROUP  II— SUB-GROUP  B 

HIS  sub-group  comprises  the  following  five  metals  : 

Beryllium  (glucinum)  (Be)  .         .        .         .  9.1 

Magnesium  (Mg)                  .        .        •        .  24.3 

Zinc  (Zn) 65.4 

Cadmium  (Cd) 112.4 

Mercury  (Hg)       .                 ....  200.0 

Corresponding  with  their  position  in  the  periodic  table,  they  function 
as  divalent  elements  only  (regarding  mercurous  compounds  cf.  p.  472). 
They  are  considerably  less  electro-positive  than  the  metals  so  far 
considered,  and  the  electro-positive  character  diminishes  from  mag- 
nesium to  mercury.  From  this  it  follows  at  once,  since  the  hydroxides 
are  relatively  weak  bases,  that  the  halogen  compounds  undergo 
hydrolysis,  and  that  there  is  considerable  tendency  to  the  forma- 
tion of  basic  salts  ;  further,  that  the  hydroxides  and  carbonates 
are  less  stable  than  those  of  the  alkalis  and  alkaline  earth  metals. 

BERYLLIUM  (GLUCINUM) 

Symbol,  Be.     Atomic  Weight=9.i. 

Occurrence — Beryllium  is  a  comparatively  rare  element.  The  chief  source 
is  beryl,  a  silicate  of  aluminium  and  beryllium,  3BeO,Al2O3,6SiO2.  Beryl 
coloured  green  by  traces  of  impurities  is  called  emerald;  when  coloured 
bluish-green,  aquamarine.  Chrysoberyl,  another  mineral  containing  beryllium, 
has  the  formula  BeO,Al2O3. 

Preparation  Of  Metal — Beryllium  is  most  readily  obtained  by  strongly  heating 
beryllium  potassium  fluoride,  BeF2'KF,  with  metallic  sodium  and  excess  of  sodium 
chloride.  After  treatment  with  water,  crystals  of  practically  pure  beryllium  are 
found  in  the  residue. 

Properties — Beryllium  is  a  silver-white  metal  of  density  1.85;  it  melts  below 
1000°.  It  is  stable  in  the  air  at  ordinary  temperature,  but  on  heating  becomes 
coated  with  a  layer  of  oxide,  which  retards  further  oxidation.  It  is  not  affected 
by  water  even  at  100°.  It  is  readily  dissolved  by  dilute  hydrochloric  and 
sulphuric  acids,  but  nitric  acid,  whether  dilute  or  concentrated,  has  very  little 

460 


MAGNESIUM  461 

action  on  it.  It  dissolves  readily  in  potassium  or  sodium  hydroxide,  with 
evolution  of  hydrogen  and  formation  of  the  compound  Be(ONa)2.  In  this 
and  other  respects  it  shows  a  remarkable  resemblance  to  aluminium. 

Compounds  of  Beryllium — Beryllium  oxide,  BeO,  is  obtained  by  heating 
the  hydroxide  to  440°;  also  by  heating  the  carbonate  or  sulphate.  It  is  a 
white  powder ;  which  under  ordinary  circumstances  is  soluble  in  dilute  acids, 
but  after  being  strongly  ignited  is  much  less  soluble. 

Beryllium  Hydroxide,  Be(OH)2,  is  obtained  as  a  gelatinous  precipitate  when 
an  alkali  hydroxide  is  added  to  the  solution  of  a  beryllium  salt.  When  heated 
at  440°,  it  loses  water  and  forms  beryllium  oxide.  It  dissolves  in  excess  of  alkali 
hydroxide,  forming  alkali  beryllate,  e.g.  Be(ONa)2,  but  is  reprecipitated  on 
prolonged  boiling. 

Beryllium  Chloride,  BeCl2,  is  obtained  in  the  anhydrous  condition  by  heating 
the  oxide,  mixed  with  charcoal,  in  a  current  of  chlorine.  It  forms  a  colourless, 
minutely  crystalline  powder,  which  fumes  in  the  air  like  phosphorus  penta- 
chloride.  It  is  readily  soluble  in  water,  and  on  evaporating  the  solution  the 
tetrahydrate,  BeCl2,4H2O,  separates  in  monoclinic,  deliquescent  crystals.  This 
compound  cannot  be  completely  dehydrated  on  heating,  as  under  these 
circumstances  the  water  reacts  with  the  salt,  hydrogen  chloride  escapes,  and 
the  oxide  remains  behind : 

BeCl2+  H2O_>BeO+2HCl. 

Beryllium  Carbonate — As  Beryllium  is  a  weak  base  and  carbonic  acid  is 
a  weak  acid,  it  is  not  surprising  that  owing  to  hydrolysis  basic  carbonates 
are  obtained  by  double  decomposition  between  a  soluble  beryllium  salt  and  a 
soluble  carbonate.  Of  these,  the  compound  BeCO3,2Be(OH)2  is  best  known. 
The  normal  carbonate,  BeCO3>4H2O,  is  obtained  by  evaporating  the  aqueous 
solution  in  an  atmosphere  of  carbon  dioxide  (cf.  magnesium  carbonates). 
The  salt  is  dehydrated  on  heating  to  100°,  and  at  a  slightly  higher  temperature 
begins  to  lose  carbon  dioxide. 

Beryllium  Sulphate,  BeSO4— This  salt  is  obtained  by  dissolving  beryllium 
oxide  in  dilute  sulphuric  acid  ;  on  concentrating  the  solution,  it  separates  as 
BeSO4,4H2O  in  large  octahedral  crystals.  It  is  very  soluble  in  water.  On 
heating  the  crystals  practically  all  the  water  escapes ;  at  a  slightly  higher 
temperature  sulphur  trioxide  is  driven  off  and  the  oxide  remains. 

MAGNESIUM 

Symbol,  Mg.     Atomic  Weight=24-32 

Occurrence — In  the  combined  state,  magnesium  is  one  of  the 
most  abundant  of  the  elements.  In  the  Stassfurt  deposits,  it 
occurs  as  chloride  in  carnallite,  MgCl2,KCl,6H2O,  and  tachydrite, 
CaCl2,2MgCl2,i2H2O,  and  as  sulphate  in  kieserite,  MgSO4,H2O  kainite, 
KCl,MgSO4,3H2O,  and  other  compounds.  As  carbonate,  it  forms 
magnetite,  MgCO3,  and  is  a  constituent  of  dolomite,  MgCO3,CaCO3, 
which  forms  whole  mountain  ranges.  As  silicate,  it  occurs  in 
talc,  Mg3H2(SiO3)4,  serpentine,  Mg3Si2O7,2H2O,  meerschaum, 


462     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Mg2Si3O8,2H2O,    asbestos    and    other    compounds.       As     chloride 
and  sulphate,  it  is  found  in  many  mineral  springs. 

Preparation  of  Metal — Magnesium  is  now  usually  pre- 
pared by  the  electrolysis  of  dehydrated  carnallite,  MgCl2,KCl. 
The  salt  is  kept  at  a  temperature  above  the  melting-point  in  an 
iron  vessel,  the  walls  of  which  serve  as  cathode,  and  the  carbon 
anode,  surrounded  by  a  porcelain  tube  to  conduct  away  the  chlorine? 
is  introduced  through  a  hole  in  the  cover. 

Formerly  magnesium  was  prepared  by  heating  the  double  chloride 
of  magnesium  and  sodium,  MgCl2'NaCl(p.463),  with  metallic  sodium. 
A  convenient  modification  of  this  method,  still  in  use,  is  to  treat 
dehydrated  carnallite  with  metallic  sodium. 

Properties — Magnesium  is  a  silvery-white,  fairly  hard  metal  of 
density  1.75  ;  it  melts  about  600°  and  boils  about  1100°.  It  is  fairly 
malleable;  on  heating  it  becomes  ductile  and  can  be  drawn  into 
wire.  It  is  stable  in  dry  air,  but  in  moist  air  becomes  covered  with 
a  coating  of  oxide.  When  heated  to  redness  in  air,  it  burns  with  a 
brilliant  white  flame,  the  high  luminosity  of  which  is  due  to  the 
great  amount  of  heat  given  out  and  to  the  non-volatility  of  the 
oxide  formed.  For  this  reason  it  is  used  in  signalling  and  in 
pyrotechny.  Further,  as  the  flame  is  very  rich  in  chemically 
active  rays,  it  is  used  as  a  flash-light  in  photography ;  for  this 
purpose  magnesium  powder  is  blown  into  a  spirit  flame.  It  acts 
on  water  very  slowly  even  at  100°,  but  when  steam  is  passed  over 
heated  magnesium,  the  metal  burns  vigorously  to  the  oxide  and 
hydrogen  is  set  free.  It  is  readily  dissolved  by  acids  with  evolu- 
tion of  hydrogen.  It  is  interesting  to  note  that  magnesium  liberates 
hydrogen  from  nitric  acid,  a  property  possessed  by  no  other  metal. 
Unlike  beryllium  and  zinc  it  is  not  affected  by  alkali  hydroxides, 
even  on  boiling. 

Magnesium  Oxide,  MgO,  is  formed  by  burning  the  metal 
in  air,  or  by  heating  the  hydroxide  or  carbonate.  It  is  a  white, 
very  light  powder,  which  is  used  in  medicine  under  the  name  of 
calcined  magnesia.  It  is  a  very  infusible  substance  (it  can,  how- 
ever, be  fused  in  the  electric  furnace),  and  is  therefore  used  in 
making  fire-brick,  crucibles,  etc.  It  combines  readily  with  water 
to  form  magnesium  hydroxide,  Mg(OH)2. 

Magnesium  Hydroxide,  Mg(OH)2,  is  obtained  by  the  action 
of  water  on  the  oxide,  or  by  adding  sodium  or  potassium  hydroxide 
to  the  solution  of  a  magnesium  salt.  It  is  a  white  powder  which 
loses  water  at  a  red  heat,  forming  the  oxide.  It  is  very  slightly 


MAGNESIUM  463 

soluble   in  water,  but   sufficient  is  taken  up   to  make  the  solution 

alkaline  to  litmus.     It  is  fairly  soluble  in  solutions  of  ammonium 

salts,  e.g.  ammonium  chloride.  The   reaction   in   this  case  is   re- 
presented by  the  equation 


and  depends  upon  the  fact  that  the  OH'  ion  concentration  in 
ammonium  hydroxide  solution  is  very  small,  especially  in  the 
presence  of  ammonium  salts.  Magnesium  hydroxide  can  only 
remain  undissolved  when  the  solubility  product  of  Mg"  and  OH' 
ions  is  exceeded.  In  the  above  instance,  OH'  ions  are  taken  up 
to  form  practically  non-ionised  ammonium  hydroxide,  as  repre- 
sented by  the  upper  arrow,  and  if  sufficient  ammonium  chloride 
is  added  the  OH'  ion  concentration  is  reduced  to  such  an  extent 
that  the  solubility  product  is  no  longer  reached  ;  in  other  words, 
magnesium  hydroxide  is  dissolved. 

A  paste  of  magnesium  hydroxide  and  water  absorbs  carbon  dioxide 
from  the  air,  and  sets  into  a  hard  mass  of  carbonate  ;  it  is  therefore 
a  useful  cement. 

Magnesium  Chloride,  MgCl2,  is  obtained  by  dissolving 
magnesium,  the  oxide  or  carbonate,  in  hydrochloric  acid;  on  con- 
centrating the  solution  it  separates  in  monoclinic  crystals  as 
MgCl2,6H2O.  Commercially  it  is  obtained  at  Stassfurt  from  carnallite, 
MgCl2,KCl,6H2O.  As  potassium  chloride  is  the  less  soluble  of  the 
two  salts,  the  greater  part  of  it  is  first  removed  by  crystallization, 
and  from  the  mother  liquor  MgCl2,6H2O  is  obtained  on  further  con- 
centration. The  crystals  are  deliquescent. 

The  hexahydrate  cannot  be  completely  dehydrated  by  heat,  as  the 
chloride  undergoes  partial  hydrolysis,  the  oxide  being  formed  and 
hydrogen  chloride  given  off: 


The  anhydrous  salt  is  obtained  by  heating  the  metal  in  a  current  of 
chlorine,  or  from  magnesium  ammonium  chloride,  MgCl2,NH4Cl,6H2O. 
When  the  latter  compound  is  heated,  it  first  loses  water,  then  at  a 
higher  temperature  ammonium  chloride,  and  finally  the  anhydrous 
salt  remains.  Anhydrous  magnesium  chloride  occurs  in  lustrous 
crystalline  leaflets  ;  it  melts  at  708°  and  can  be  distilled  unchanged 
in  dry  air,  but  is  decomposed  on  heating  in  moist  air,  in  accordance 
with  the  above  equation. 


464     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

A  number  of  basic  chlorides  of  magnesium  have  been  described, 
but  their  formulae  have  not  been  definitely  established. 

Magnesium  Sulphate  (Epsom  salts),"MgSO4,  occurs  in 
Stassfurt  deposits  as  kieserite,  MgSO4,H2O,  and  as  a  constituent  of 
a  number  of  double  salts  (p.  415).  It  is  also  found  in  many  mineral 
springs,  notably  in  that  at  Epsom,  from  which  it  was  first  obtained. 
It  is  usually  obtained  commercially  by  treating  kieserite  (a  com- 
paratively insoluble  salt)  with  hot  water,  and  setting  the  solution 
aside  to  crystallize. 

Magnesium  sulphate  is  usually  met  with  as  the  heptahydrate, 
MgSO4,7H2O,  in  rhombic  prisms.  Another  unstable  form  of  the 
heptahydrate  (monoclinic  prisms)  sometimes  separates  from  super- 
saturated solutions  at  room  temperature.  At  60  to  70°  the  hexa- 
hydrate  separates  from  solution  in  monoclinic  crystals,  and  above  68° 
kieserite,  MgSO4,H2O,  is  obtained.  Other  hydrates, e.g.  MgSO4,i  2 H2O, 
are  also  known.  All  the  higher  hydrates,  when  heated  at  150°,  lose 
water  and  form  the  monohydrate ;  above  200°  the  anhydrous  salt  is 
obtained.  At  o°  100  grams  of  water  dissolve  26.9  grams,  at  15°  33.8 
grams,  and  at  30°  40.9  grams  of  MgSO4. 

An  acid  sulphate,  Mg(HSO4)2,has  been  obtained  in  prismatic  crystals 
from  a  hot  solution  of  magnesium  sulphate  in  sulphuric  acid. 

Double  salts  of  magnesium  sulphate  with  the  alkali  sulphates,  of 
the  type  MgSO4,M2iSO4  are  known.  The  potassium  and  ammonium 
salts  crystallize  with  6H2O,  the  sodium  salt  with  4H2O.  The  potassium 
salt  occurs  naturally  as  schonite. 

Magnesium  Ammonium  Phosphate,  MgNH4PO4,6H2O, 
is  always  precipitated,  on  account  of  its  very  slight  solubility  in  water, 
when  Mg",  NH4'  and  PO4"'  ions  are  brought  together,  e.g.  by  adding 
ammonium  chloride,  ammonia  and  sodium  phosphate  to  a  solution 
of  magnesium  sulphate.  Although  only  slightly  soluble  in  water,  it  is 
still  less  soluble  in  solutions  containing  NH4'  ions  and  free  ammonia. 
It  is  used  in  the  quantitative  estimation  of  magnesium  salts  and 
of  phosphates.  On  strongly  heating,  magnesium  pyrophosphate, 
Mg2P2O7,  is  obtained: 

2MgNH4P04->Mg2P207  +  2NH3+H20. 

Magnesium  Carbonates  —  The  normal  carbonate  occurs 
naturally  as  magnesite,  MgCO3,  in  rhombohedral  crystals,  and  a 
basic  carbonate,  3MgCO3,Mg(OH)2,3H2O,  as  hydromagnesile,  in 
monoclinic  crystals.  As  both  base  and  acid  are  weak,  only  basic 
carbonates  are  obtained  by  double  decomposition  between  mag- 


ZINC  465 

nesium  salts  and  carbonates  in  solution.  The  composition  of 
these  basic  carbonates  depends  on  the  conditions  of  precipitation. 
When  solutions  of  magnesium  sulphate  and  sodium  carbonate  are 
mixed  at  room  temperature,  the  precipitate  has  the  approximate  com- 
position 3MgCO3,Mg(OH)2,3H2O,  the  same  as  hydromagnesite.  The 
carbonate  thus  obtained  is  known  as  light  magnesium  carbonate.  A 
much  denser  precipitate,  which  may  be  rather  more  basic,  is  obtained 
by  carrying  out  the  precipitation  in  concentrated  solution  at  the 
boiling-point ;  it  is  known  as  heavy  magnesium  carbonate.  There  is 
no  evidence  that  these  basic  carbonates  are  other  than  mixtures  of 
carbonate  and  hydroxide,  and  in  fact  the  composition  of  the  com- 
mercial article  varies  within  wide  limits. 

When  the  basic  carbonate  is  suspended  in  water  and  carbon  dioxide 
passed  through  it,  the  carbonate  dissolves,  and  the  solution  doubtless 
contains  the  bicarbonate,  Mg(HCO3)2.  From  this  solution  below 
1 6°  the  normal  carbonate  separates  as  pentahydrate,  and  above 
16°  as  trihydrate,  MgCO3,3H2O,  in  colourless  crystals.  The  an- 
hydrous carbonate  can  be  obtained  by  heating  a  solution  of  the 
carbonate  in  excess  of  carbon  dioxide  to  150°  in  a  vessel  provided 
with  a  porous  stopper,  so  that  the  carbon  dioxide  can  escape  slowly. 

On  heating  at  200°  both  normal  and  basic  carbonates  are  com- 
pletely decomposed,  with  formation  of  the  oxide. 

Tests  for  Magnesium— From  solutions  of  magnesium  salts, 
the  alkali  hydroxides  precipitate  magnesium  hydroxide,  insoluble  in 
excess  of  alkali.  For  the  reasons  already  stated,  ammonia  does  not 
precipitate  magnesium  hydroxide  from  solutions  containing  ammonium 
chloride.  From  solutions  of  magnesium  salts  containing  ammonium 
chloride  and  excess  of  ammonia,  sodium  phosphate  throws  down  a 
white  crystalline  precipitate  of  magnesium  ammonium  phosphate, 
MgNH4PO4. 

ZINC 

Symbol,  Zn.     Atomic  weight=65.37.     Molecular  weight=65.37. 

Occurrence — Zinc  is  fairly  abundant  in  nature  as  calamine  or 
smithsonite,  ZnCO3,  zinc  blende,  ZnS,franklinite,  Zn(FeO2)2,  and  red 
zinc  ore,  ZnO.  The  majority  of  the  zinc  ores  contain  a  little  cadmium. 

Preparation  of  Metal— The  first  stage  in  obtaining  the  metal 

is  to  convert  the  ores  to  oxide.     For  this  purpose  the  sulphide  is 

roasted  in  air  ;  the  carbonate  only  requires  to  be  heated.     The  oxide 

is  then  mixed  with  powdered  coal   or  coke  and  heated  strongly  in 

3° 


466     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

earthenware  retorts,  when  zinc  distils  over  and  is  collected  in  iron  or 
fireclay  receivers.  The  portion  which  first  passes  over  is  condensed  in 
the  cold  receiver  as  a  fine  powder  termed  zinc  dust.  Besides  the  metal, 
zinc  dust  contains  a  few  per  cent,  of  oxide,  and  also  some  cadmium 
which,  being  more  volatile  than  zinc,  passes  over  first.  As  the  retorts 
heat  up,  the  zinc  melts,  and  can  be  drawn  off  and  cast  into  sticks. 

Zinc  thus  prepared  is  very  impure,  containing  lead,  iron,  cadmium, 
arsenic  and  other  elements.  It  can  be  partially  purified  by  redis- 
tillation. Perfectly  pure  zinc  is  obtained  by  dissolving  the  metal  in 
acid,  and  precipitating  the  carbonate,  which  is  then  heated  and  the 
oxide  reduced  with  pure  carbon. 

Properties— Zincl is  a  bluish-white  metal  of  density  6.9  to  7.2  ;  it 
melts  at  419°  and  boils  at  920°.  It  is  usually  obtained  in  hexagonal 
crystals.  Ordinary  zinc  is  moderately  brittle  at  room  temperature, 
at  150°  it  is  malleable  and  ductile,  at  200°  it  has  again  become  so 
brittle  that  it  can  be  powdered  in  a  mortar.  This  behaviour  is 
doubtless  connected  with  changes  in  the  nature  of  the  crystals.  The 
vapour  density  of  zinc  is  about  32.7,  from  which  it  follows  that  it  is 
monatomic  in  the  state  of  vapour. 

In  the  air,  zinc  becomes  coated  with  an  extremely  thin  film  of  oxide 
(or  perhaps  basic  carbonate)  which  protects  it  against  further  oxida- 
tion. On  heating  in  air,  it  burns  with  a  bluish  flame,  forming  the 
oxide.  It  does  not  decompose  water  at  100°.  The  commercial  metal 
dissolves  readily  in  dilute  acids,  hydrogen  being  given  off,  but  very 
pure  zinc  is  scarcely  affected  by  acids.  This  difference  in  behaviour 
of  pure  and  impure  zinc  is  not  entirely  understood.  In  the  latter  case, 
differences  of  potential  are  probably  set  up  between  zinc  and  the 
impurities  (iron,  lead,  etc.)  and  the  hydrogen  is  liberated  at  these 
metals  as  in  an  ordinary  battery  (cf.  p.  432).  With  pure  zinc,  on  the 
other  hand,  the  hydrogen  probably  accumulates  on  the  surface  and 
sets  up  a  contrary  E.M.F.  which  brings  the  reaction  to  a  standstill. 

When  boiled  with  solutions  of  alkali  hydroxides,  zinc  dissolves  with 
liberation  of  hydrogen  and  formation  of  alkali  zincate  : 

Zn  +  2KOH->Zn(OK)2  +  H2  f  . 

Uses  of  Zinc — Zinc  is  largely  used  in  preparing  galvanized  iron, 
which  is  iron  coated  with  zinc  in  order  to  protect  it  against  oxidation. 
Galvanized  iron  is  not  prepared  by  electrolytic  deposition,  as  the  name 
appears  to  imply,  but  by  dipping  clean  iron  sheets  into  molten  zinc. 
Zinc  in  the  form  of  sheets  is  also  sometimes  used  for  covering 
roofs,  etc. 


ZINC  467 

Zinc  is  an  important  constituent  of  certain  alloys,  such  as  brass  and 
German  silver  (p.  427),  and  is  present  in  small  proportion  in  copper 
coins.  With  some  metals,  e.g.  cadmium,  it  is  completely  miscible  in 
the  fused  state  ;  with  others,  such  as  lead,  it  is  only  slightly  miscible 
(ff.  p.  436). 

Zinc  Oxide,  ZnO,  is  obtained  by  burning  zinc  in  the  air  or  by 
strongly  heating  the  basic  carbonate.  It  forms  a  white  amorphous 
powder,  which  becomes  yellow  on  heating  but  regains  its  original 
colour  on  cooling.  It  does  not  combine  with  water  to  form  the 
hydroxide,  but  dissolves  readily  in  acids  to  form  the  corresponding 
salts.  It  is  used  as  a  pigment  under  the  name  of  zinc-white.  It 
has  one  advantage  over  white  lead  in  not  being  blackened  by 
hydrogen  sulphide. 

Like  magnesium  oxide,  it  becomes  incandescent  on  heating,  and 
even  at  a  white  heat  does  npt  fuse  or  decompose. 

Zinc  Hydroxide,  Zn(OH)2,  is  formed  as  a  white  gelatinous  pre- 
cipitate when  potassium,  sodium  or  ammonium  hydroxide  is  added  to 
a  solution  of  a  zinc  salt.  It  is  soluble  in  excess  of  the  alkalis  as  well 
as  of  ammonia.  As  regards  the  alkalis,  this  is  due  to  the  formation 
of  alkali  zincate,  e.g.  Na2ZnO2  ;  in  the  case  of  ammonia  to  the  forma- 
tion of  complex  cations,  Zn(NH3)4",  the  solubility  product  of  these 
and  OH7  ions  being  much  larger  than  that  of  Zn"  and  OH'  ions. 

The  weakly  basic  character  of  zinc  hydroxide  is  shown  by  the 
formation  of  complex  ions  with  ammonia  and  by  the  hydrolysis 
of  its  salts  ;  its  slightly  acidic  character  by  the  existence  of  alkali 
zincates. 

Zinc  Chloride,  ZnCl2,  is  obtained  in  the  anhydrous  form  by 
heating  zinc  in  a  current  of  chlorine,  and  in  solution  by  dissolving 
the  metal,  oxide  or  carbonate  in  hydrochloric  acid.  On  evaporating 
the  solution  partial  hydrolysis  takes  place,  and  hydrogen  chloride  is 
given  off;  the  product  contains  a  certain  proportion  of  oxychloride 

/OH 
Zn<^         ,  and  perhaps  zinc  oxide  as  well  : 

+  HOH-»Zn(OH)Cl  +  HCl. 


The  anhydrous  salt  is  usually  made  into  sticks.  It  is  very  deliquescent, 
and  is  sometimes  used  as  a  dehydrating  agent.  When  it  is  treated 
with  water,  the  oxychloride  generally  present  remains  insoluble,  but 
a  clear  solution  can  be  obtained  by  adding  hydrochloric  acid.  The 
pure  anhydrous  salt  melts  about  290°,  and  boils  at  730°.  When  zinc 
oxide  is  added  to  a  very  concentrated  solution  of  zinc  chloride,  the 


468     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

mixture  rapidly  sets  to  a  hard  mass  (mainly  owing  to  the  formation  of 
oxychloride).  * 

Zinc  chloride  is  used  for  preserving  railway  sleepers  against  decay  ; 
also  for  cleaning  metal  surfaces  in  soldering,  and  as  a  caustic. 

Zinc  Carbonates  —  The  normal  carbonate  occurs  naturally  as 
calamine  in  rhombohedral  crystals  isomorphous  with  calcspar.  It  is 
obtained  by  adding  sodium  bicarbonate  to  a  solution  of  zinc  sulphate. 

When  the  normal  alkali  carbonates  are  used,  basic  salts  are  ob- 
tained, the  composition  of  which  depends  on  the  temperature,  concen- 
tration and  other  conditions  (cf.  magnesium  carbonates).  Commercial 
zinc  carbonate  approximates  to  the  composition  ZnCO3,2Zn(OH)2. 

Zinc  Sulphate,  ZnSO4,  is  obtained  by  dissolving  zinc  in  sul- 
phuric acid.  With  the  dilute  acid  the  sulphate  and  hydrogen  are 
formed  ;  with  the  concentrated  acid  the  sulphate  is  formed  and 
sulphur  dioxide  is  given  off: 


On  the  commercial  scale,  it  is  prepared  by  roasting  the  sulphide  to 
sulphate,  and  dissolving  out  the  latter  with  water. 

Zinc  sulphate  is  usually  met  with  as  the  heptahydrate,  ZnSO4,7H2O 
(hexagonal  crystals,  isomorphous  with  MgSO4,  '/H2O),  which  separates 
from  aqueous  solution  below  39°.  The  crystals  effloresce  slightly  in 
the  air,  above  100°  they  lose  water  and  form  the  monohydrate, 
ZnSO4,H2O,  and  above  240°  the  anhydrous  salt  is  obtained.  At  o° 
loo  grams  of  water  dissolve  41.9  grams,  at  15°  50.8  grams,  and  at  35° 
66.6  grams  of  the  anhydrous  salt. 

Above  39°  the  compound  ZnSO456H2O  separates  in  monoclinic 
crystals,  isomorphous  with  MgSO4,6H2O. 

Zinc  sulphate  forms  double  salts  with  the  alkali  sulphates,  e.g. 
ZnSO4,K2SO4,6H2O,  isomorphous  with  the  corresponding  magnesium 
compounds  (p.  464). 

Zinc  Sulphide,  ZnS,  occurs  naturally  in  regular  crystals  as  zinc 
blende  (coloured  brown  or  black  by  traces  of  impurities)  and  in 
hexagonal  crystals  as  wurtzite.  It  is  obtained  as  a  white  precipitate 
when  ammonium  sulphide  is  added  to  the  solution  of  a  zinc  salt,  also 
when  hydrogen  sulphide  is  passed  into  an  alkaline  solution  of  a  zinc 
salt,  but  not  when  the  gas  is  passed  into  an  acid  solution.  In  alkaline 
solution,  owing  to  the  presence  of  a  considerable  concentration  ofZn" 
and  S"  ions,  the  solubility  product  of  ZnS  is  exceeded  and  precipita- 
tion occurs,  but  in  acid  solution,  owing  to  the  high  H'  ion  concentra- 
tion, the  S"  ion  concentration  is  reduced  to  such  an  extent  that  no 


CADMIUM  469 

precipitation  occurs.     In  neutral  solution,  owing  to  the  accumulation 
of  H'  ions  during  the  reaction,  precipitation  is  only  partial  : 


2H'-f  SO4", 

but  the  addition  of  sodium  acetate  to  the  neutral  solution  prevents  the 
accumulation  of  H*  ions  owing  to  the  small  ionisation  of  acetic  acid: 

H-  +  CH3COO'-r-Na-->CH3COOH  +  Na'. 

The  fact  that  zinc  sulphide  is  insoluble  in  acetic,  but  soluble  in  hydro- 
chloric acid,  can  be  accounted  for  on  the  same  lines. 

Although  zinc  sulphide  is  practically  insoluble  in  water,  it  has  con- 
siderable tendency  to  pass  into  colloidal  solution.  This  is  often 
observed  when  the  freshly  precipitated  compound  is  being  washed 
with  water. 

Tests  for  Zinc  —  Zinc  salts  give  no  definite  colour  to  the  Bunsen 
flame.  The  more  important  tests  are  the  precipitation,  by  means  of 
ammonium  sulphide,  of  white  zinc  sulphide,  which,  for  the  reasons 
explained  above,  is  insoluble  in  acetic  acid  but  soluble  in  hydrochloric 
acid.  The  formation  of  a  white  precipitate  on  addition  of  ammonium 
hydroxide,  soluble  in  excess  of  the  latter,  is  also  a  useful  test. 

CADMIUM 

Symbol,  Cd.     Atomic  weight  —  112.4.     Molecular  weight  =  112.  4. 

Occurrence  —  Cadmium  is  a  comparatively  rare  element.  It 
occurs  in  nature  as  greenockite,  CdS,  a  very  scarce  mineral,  and  is 
present,  generally  not  in  excess  of  I  per  cent.,  in  most  zinc  ores,  from 
which  it  is  exclusively  obtained  commercially. 

Preparation  of  Metal  —  Cadmium,  being  more  volatile  than 
zinc,  passes  over  first  when  the  oxides  are  heated  with  charcoal  in  the 
commercial  preparation  of  zinc  (p.  465).  It  can  be  separated  from 
zinc  by  fractional  distillation  or  by  precipitation  as  sulphide,  which, 
unlike  zinc  sulphide,  is  insoluble  in  dilute  hydrochloric  acid.  The 
sulphide  is  then  dissolved  in  concentrated  hydrochloric  acid,  precipi- 
tated as  carbonate,  and  the  latter  converted  by  heating  into  the 
oxide,  which  is  finally  distilled  with  carbon. 

Properties  —  Cadmium  is  a  silvery-white  crystalline  metal  of 
density  8.6  ;  it  melts  at  321°,  and  boils  at  780°.  It  is  malleable  and 
ductile  at  room  temperature.  In  the  air  at  room  temperature  it 
becomes  coated  with  a  thin  film  of  oxide  ;  on  heating  it  burns  readily 
with  formation  of  the  oxide.  It  dissolves  in  dilute  acids  with  forma- 
tion of  the  corresponding  salts. 


470    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Cadmium  Oxide,  CdO,  is  obtained  as  a  brown  powder  when 
cadmium  is  burned  in  air,  and  also  by  heating  the  carbonate.  At  a 
bright  red  heat  in  an  atmosphere  of  oxygen  it  changes  to  dark-red  cubic 
crystals.  It  is  readily  reduced  by  heating  in  a  stream  of  hydrogen. 

Cadmium  Hydroxide,  Cd(OH)2,  is  obtained  by  the  action  of 
water  on  the  oxide,  or  by  double  decomposition,  as  a  white  powder. 
It  is  very  slightly  soluble  in  water.  Unlike  zinc  hydroxide,  it  is  not 
soluble  in  sodium  or  potassium  hydroxide,  but  is  soluble  in  ammonia 
on  account  of  the  formation  of  complex  ions. 

Cadmium  Chloride,  CdCl2,  prepared  by  the  general  methods, 
separates  from  solution  in  colourless,  monoclinic  crystals  as 
CdCl2,2H2O.  On  heating  it  loses  all  its  water  without  further  decom- 
position. Cadmium  iodide,  CdI2,  occurs  in  anhydrous  crystals, 
very  soluble  in  water. 

The  behaviour  of  the  cadmium  halides  in  solution  shows  that  the 
normal  ionisation  is  small,  and  that  complex  anions  containing  the 
metal  are  present.  Aqueous  solutions  of  the  iodide  contain  a  con- 
siderable proportion  of  CdI3'  ions.  The  chloride  has  much  less 
tendency  to  complex  formation  than  the  iodide. 

Cadmium  Sulphate,  CdSO4,  obtained  by  the  general  methods 
for  a  soluble  salt,  usually  separates  from  solution  in  colourless 
crystals  as  the  hydrate,  3CdSO4,8H2O.  The  composition  of  the 
hydrate  is  rather  remarkable,  but  appears  to  be  definitely  established. 
The  crystals  effloresce  on  exposure  to  air. 

Cadmium  Sulphide,  CdS,  is  obtained  as  a  bright  yellow 
precipitate  by  passing  hydrogen  sulphide  into  the  solution  of  any 
cadmium  salt.  It  is  soluble  in  concentrated  hydrochloric  acid,  but 
insoluble  in  ammonium  sulphide.  It  is  used  as  a  pigment  in  painting. 

MERCURY 

Symbol,  Hg.     Atomic  weight =200.0.     Molecular  weight =200.0. 

Occurrence — Mercury  is  sometimes  found  in  the  free  condition 
as  small  globules  in  cavities  of  rocks,  but  the  chief  source  is  cinnabar, 
HgS,  which  forms  dark  red,  often  crystalline  masses.  This  ore  is 
found  chiefly  at  Idria  in  Austria,  Almaden  in  Spain,  California,  Peru, 
China,  and  Japan. 

Preparation  of  Metal — As  mercury  has  very  little  affinity 
for  oxygen,  it  is  readily  obtained  from  its  ores.  According  to  one 
method,  the  sulphide  is  simply  roasted  in  air  and  the  metal  condensed  : 


MERCURY 


47T 


FIG.  86. 


otherwise  it  is  heated  with  lime,  whereby  the  sulphur  is  retained  as 

calcium  sulphide  and  the  vapours  of  mercury  are  condensed.     At 

Idria   the   condensation    is   effected   in 

large  chambers  connected  in  series ;  in 

Almaden    aludels   are    used    (Fig.    86). 

These  consist   of  pear-shaped  earthen 

vessels  arranged  in  series,  the  narrow  neck  of  one  being  inserted  in 

the  base  of  the  next,  as  shown. 

The  mercury  thus  obtained  is  very  impure,  containing  dissolved 
metals,  such  as  lead,  copper,  zinc,  and  arsenic,  and  also  mechanically 
mixed  impurities.  The  latter  are  removed  by  filtering  through 
chamois  leather,  the  former  by  prolonged  shaking  with  dilute  nitric 
acid  or  with  dilute  sulphuric  acid  containing  a  little  potassium 
bichromate.  The  agitation  is  conveniently  effected  by  drawing  a 
current  of  air  through  the  liquid.  These  methods  depend  upon  the 
fact  that  the  usual  impurities  are  much  more  easily  oxidized  and 
dissolved  than  mercury  itself.  The  treatment  with  nitric  acid  is 
conveniently  carried  out  in  the  apparatus  represented  in  Fig.  87. 
The  funnel  A  ends  in  a  capillary,  so  that  the  mercury  enters  the  acid 
(which  is  contained  in  a  long  glass  tube,  B)  in  a  fine  stream.  It  is 
then  washed  and  dried.  If  required  perfectly 
pure,  it  is  finally  distilled  in  a  vacuum. 

Properties — Mercury  is  the  only  metal  which 
is  liquid  at  room  temperature.  To  this  property 
it  owes  the  names  quicksilver  and  hydrargyrum 
(from  #5w/>,  water,  and  apyvpos,  silver).  The  solid 
metal,  which  forms  octahedral  crystals,  and  is  very 
malleable  and  ductile,  melts  at  -39°,  and  the 
liquid  boils  at  357°.  The  density  at  o°  is  13.596, 
at  20°  13.546.  Mercury  is  slightly  volatile  even 
at  room  temperature,  as  is  shown  by  the  fact  that 
the  surface  of  a  gold  leaf  suspended  over  the  liquid 
in  a  closed  space  becomes  amalgamated.  The 
vapour  of  mercury  conducts  the  electric  current 
at  high  temperatures  and  gives  a  peculiar  greenish 
light.  This  is  the  principle  of  the  "  mercury  vapour 
lamps  "  now  widely  used. 

Mercury  is  not  affected  by  oxygen  at  room  tem- 
FIG.  87.  perature,  but  on  prolonged  heating  at  its  boiling- 

point  in  contact  with  air   red  mercuric  oxide  is 
formed  (p.  27).     Hydrochloric  acid  and  dilute  sulphuric  acid  have 


472     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

no  action  on  mercury,  but  when  the  metal  is  heated  with  concen- 
trated sulphuric  acid  mercuric  sulphate  is  formed  and  sulphur  dioxide 
given  off: 

Hg+2H2SO4-»HgSO4+SO2+2H2O. 

Concentrated  nitric  acid  acts  on  mercury  with  formation  of  mercuric 
nitrate,  whilst  with  moderately  dilute  acid  and  excess  of  mercury 
mercurous  nitrate,  Hg2(NO3)2,  is  the  chief  product. 

Amalgams — Alloys  of  mercury  with  other  metals  are  called 
amalgams.  They  are  usually  prepared  by  the  direct  addition  of  the 
other  metal  to  mercury,  sometimes  also  by  electrolytic  liberation  of 
the  metal  at  a  mercury  cathode.  Amalgams  may  be  liquid  or  solid, 
depending  upon  the  nature  and  proportion  of  the  second  metal. 
Nearly  all  the  metals  form  amalgams  with  mercury,  platinum  and 
iron  being  almost  the  only  exceptions.  In  the  majority  of  cases  the 
amalgams  contain  definite  chemical  compounds,  but  in  some  instances 
the  metals  simply  mix  without  combining. 

The  alkali  metals  in  certain  proportions  form  solid  amalgams  with 
mercury,  and  a  number  of  definite  compounds  have  been  isolated, 
e.g.  NaHg6,  NaHg6,  LiHg6,  KHg12,  etc.  The  use  of  sodium  amalgam 
in  preparing  hydrogen  has  already  been  referred  to.  The  amalgams 
of  zinc  and  of  cadmium  are  of  great  importance  in  connexion  with 
cells  used  as  standards  of  E.M.F.  The  use  of  mercury  in  extracting 
gold  from  its  ores  has  already  been  mentioned.  Chemical  com- 
pounds are  probably  formed  in  this  case,  but  their  formulae  have  not 
been  established.  Lead  readily  forms  an  amalgam  with  mercury, 
but  in  this  case  no  chemical  compounds  appear  to  be  formed. 

Compounds  of  Mercury— Mercury  forms  two  series  of  salts, 
in  both  of  which  it  is  bivalent.  In  the  mercurous  compounds  the 
bivalent  ion  contains  two  atoms  of  mercury,  e.g.  Hg2Cl2,  Hg2SO4, 
whereas  the  mercuric  compounds  contain  a  single  bivalent  atom,  e.g. 
HgCl2,  HgSO4.  The  mercurous  compounds  are  sometimes  written  as 
if  they  contained  univalent  mercury,  e.g.  HgCl ;  but  for  reasons  which 
cannot  be  considered  here  the  double  formula  is  regarded  as  the 
correct  one. 

Both  mercurous  and  mercuric  oxide  are  weak  bases,  so  that  many 
of  the  salts  are  hydrolyzed  in  solution.  On  the  other  hand,  mercury 
differs  from  copper  and  other  weak  bases  in  that  it  does  not  combine 
with  ammonia  to  form  complex  ions  ;  but  reaction  takes  place  in  a 
different  way. 

As  the  terms  imply,  mercuric  salts  can  in  general  be  obtained  from 


MERCURY  473 

mercurous  salts  by  oxidation,  and  mercurous  salts  from  mercuric 
salts  by  reduction.  When  excess  of  mercury  is  used,  the  mercurous 
salts  are  generally  obtained.  The  soluble  mercuric  salts  are  very 
poisonous. 

MERCUROUS  SALTS 

Mercurous  Oxide,  Hg2O,  is  obtained  by  the  action  of  alkali 
on  mercurous  chloride  : 


It  is  a  dark  brown  powder,  which  slowly  decomposes  at  room  tempe- 
rature, rapidly  on  exposure  to  light,  into  mercuric  oxide  and  mercury. 
The  corresponding  hydroxide,  Hg2(OH)2,  has  not  been  obtained  ;  it 
is  presumably  highly  unstable. 

Mercurous  Chloride  (calomel),  Hg2Cl2,  is  obtained  by  add- 
ing a  chloride  to  a  solution  of  mercurous  nitrate,  or  more  usually  by 
heating  a  mixture  of  mercurous  sulphate  (or  mercuric  sulphate  and 
mercury)  with  sodium  chloride  : 


If  the  vapour  is  condensed  in  a  large  chamber  it  is  obtained  as  a  fine 
powder  ;  if  in  a  small  chamber  it  forms  a  solid  mass.  Calomel  is 
also  formed  by  heating  a  mixture  of  mercuric  chloride  and  mercury. 

Mercurous  chloride  usually  occurs  as  a  white  powder,  which  darkens 
on  exposure  to  light  owing  to  the  formation  of  mercury  and  mercuric 
chloride  : 

Hg2Cl2:±Hg+HgCl2. 

Its  vapour  density  above  400°  is  about  118,  which  appears  to  indicate 
that  the  formula  is  HgCl.  It  has  been  found,  however,  that  the 
vapour  rapidly  amalgamates  gold  leaf,  so  that  free  mercury  must  be 
present,  and  this  is  readily  accounted  for  if  the  calomel  dissociates  on 
heating  into  mercuric  chloride  and  mercury,  as  in  the  above  equation. 
Baker  has  found  that  in  the  perfectly  dry  condition  the  density  corre- 
sponds with  the  formula  Hg2Cl2,  an  observation  which  supports  the 
view  that  the  low  density  of  the  vapour  in  the  ordinary  condition  is 
due  to  dissociation.  Concentrated  solutions  of  hydrochloric  acid  and 
other  chlorides  dissolve  calomel  to  some  extent,  the  above  equilibrium 
being  displaced  towards  the  right. 

Calomel  is  largely  used  in  medicine. 

Mercurous  Iodide,  Hg2I2,  is  obtained  by  direct  combination 
of  its  elements.  It  is  a  yellow  to  yellowish-green  powder,  and  is 
unstable,  tending  to  break  up  into  mercuric  iodide  and  mercury. 


474     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Mercurous  Nitrate,  Hg2(NO3)2,  is  obtained  by  the  action  of 
dilute  nitric  acid  (the  concentrated  acid  diluted  with  two  or  three  parts 
of  water)  on  excess  of  mercury  at  a  gentle  heat.  It  separates  from 
solution  as  Hg2(NO3)2,2H2O  in  monoclinic  crystals,  which  effloresce 
in  the  ail.  When  treated  with  excess  of  water,  basic  salts  are  preci- 
pitated, e.g.  Hg2(NO3)2,Hg2O,H2O,  an  orange-yellow  powder ;  but 
with  excess  of  nitric  acid  a  clear  solution  is  obtained. 

Mercurous  Sulphate,  Hg2SO4,  is  obtained  by  adding  a  soluble 
sulphate  to  a  solution  of  mercurous  nitrate,  or  by  rubbing  together 
mercuric  sulphate  with  the  calculated  quantity  of  mercury.  It  forms 
a  heavy  white  crystalline  powder,  very  slightly  soluble  in  water. 


MERCURIC  SALTS 

Mercuric  Oxide,  HgO,  is  obtained  as  a  red  crystalline  powder 
by  prolonged  heating  of  mercury  in  contact  with  air  at  its  boiling- 
point  in  a  closed  vessel  (p.  27)  ;  more  readily  by  strongly  heating 
mercurous  or  mercuric  nitrate.  It  is  obtained  as  a  yellow  powder 
by  double  decomposition  between  mercuric  chloride  and  sodium  or 
potassium  hydroxide.  According  to  Ostwald,  the  only  difference 
between  the  two  forms  is  the  state  of  division,  associated,  with 
a  slight  difference  in  solubility.  By  prolonged  grinding  in  a 
mortar  the  red  oxide  becomes  yellow,  and  its  solubility  in  water 
slightly  increases.  The  red  oxide  becomes  black  on  heating,  due  to 
a  new  modification  stable  at  high  temperatures,  but  returns  to  its 
original  colour  on  cooling. 

Mercuric  Chloride  (corrosive  sublimate),  HgCl2,  is  pre- 
pared on  the  large  scale  by  subliming  a  mixture  of  mercuric  sulphate 
and  sodium  chloride,  a  little  manganese  dioxide  being  added  to  prevent 
the  formation  of  calomel.  It  forms  rhombic  crystals,  which  are 
anhydrous.  At  o°  100  parts  of  water  dissolve  5.73  grams,  at  20°  7.39 
grams,  and  at  40°  9.62  grams  of  the  salt.  It  is  still  more  soluble  in 
alcohol  and  in  ether  than  in  water. 

It  is  remarkable  that  mercuric  chloride  is  only  very  slightly  ionised 
in  aqueous  solution  ;  in  this  respect  it  differs  from  all  other  metallic 
chlorides  except  that  of  cadmium.  This  perhaps  accounts  for  its 
solubility  in  many  organic  solvents,  and  also  for  the  fact  that  the  salt 
is  scarcely  affected  by  concentrated  sulphuric  acid  or  nitric  acid  ;  it 
can,  in  fact,  be  volatilized  unchanged  from  boiling  sulphuric  acid.  It 
is  only  very  slightly  hydrolyzed  in  aqueous  solution. 

With   hydrochloric    acid   mercuric  chloride  forms   a   number  of 


MERCURY  475 

crystalline  double  salts;  for  example,  HgCl2,2HCl,7H2O  and 
2HgCl2,HCl,6H2O  ;  and  with  the  alkali  chloride  it  forms  compounds 
of  similar  type,  e.g.  HgCl2,KCl,H2O  and  HgCl2.2KCl,H2O.  In 
aqueous  solutions  of  these  compounds  the  mercury  is  mainly  present 
as  a  constituent  of  complex  ions,  such  as  HgCl3'  and  HgCl4". 

Mercuric  chloride  is  a  powerful  poison.  It  is  also  largely  used  as 
an  antiseptic  and  for  preserving  anatomical  specimens,  etc.  When 
used  as  an  antiseptic  sodium  chloride  is  usually  added  to  the  solution. 
This  has  the  advantage  that  the  solutions  keep  better,  but  they  are 
also  less  active,  as  owing  to  the  formation  of  complex  ions,  which  are 
inactive,  the  Hg"  ion  concentration  is  considerably  reduced.  The 
best  antidote  for  mercuric  chloride  is  white  of  egg,  with  which  it  forms 
an  insoluble  compound. 

Mercuric  Iodide,  HgI2,  is  obtained  by  rubbing  together  the 
calculated  quantities  of  mercury  and  iodine  in  a  mortar  with  the  addi- 
tion of  a  few  drops  of  alcohol,  or  by  double  decomposition  between 
mercuric  chloride  and  potassium  iodide  in  aqueous  solution.  When 
first  precipitated  it  is  yellow,  but  rapidly  changes  to  red.  The  red 
form  (tetragonal  crystals)  is  stable  at  room  temperature,  but  on  heat- 
ing changes  to  the  yellow  modification  (rhombic  crystals).  The 
transition  temperature  red  ^t  yellow  lies  at  126°. 

Mercuric  iodide  is  very  slightly  soluble  in  water,  but  readily  dis- 
solves in  potassium  iodide  solution.  The  solution  contains  complex 
salts,  such  as  KHgI3  and  K2HgI4,  both  of  which  have  been  obtained 
in  the  solid  form.  A  solution  of  mercuric  iodide  in  potassium 
iodide  solution  with  excess  of  alkali  is  known  as  Nessler's  reagent, 
and  is  used  for  the  detection  and  estimation  of  small  quantities  of 
ammonia  (p.  477). 

Mercuric  Cyanide,  Hg(CN)2,  obtained  by  dissolving  mercuric 
oxide  in  hydrocyanic  acid,  occurs  in  colourless  quadratic  crystals, 
readily  soluble  in  water.  The  salt  is  scarcely  ionised  at  all  in  aqueous 
solution,  and  in  accordance  with  this  gives  no  precipitate  with  alkali 
or  with  potassium  iodide,  but  gives  a  black  precipitate  with  hydrogen 
sulphide  This  behaviour  can  readily  be  accounted  for  on  the  theory 
of  the  solubility  product  (p.  439). 

Mercuric  Nitrate,  Hg(NO)2,  is  prepared  by  boiling  mercury 
with  excess  of  nitric  acid  for  some  time.  When  the  solution  ievapo- 
rated  over  concentrated  sulphuric  acid  large  colourless  deliquescent 
crystals  of  the  formula  2HgNO3'H2O  separate.  When  mercuric 
nitrate  is  treated  with  water,  basic  salts,  including  the  compound 
Hg(NO3)2,2HgO,  are  formed  by  hydrolysis.  If  excess  of  water  is 


476     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

added  and  the  mixture  heated,  complete  hydrolysis  into  mercuric 
oxide  and  nitric  acid  occurs. 

Mercuric  Sulphide,  HgS,  occurs  naturally  as  cinnabar  (red 
crystals).  It  is  obtained  in  black,  amorphous  form  by  rubbing 
together  mercury  and  sulphur,  and  also  by  passing  hydrogen  sul- 
phide into  the  neutral  or  acid  solution  of  any  mercuric  salt.  The 
red  crystalline  stable  modification  is  obtained  by  subliming  the  black 
form,  or  by  warming  the  latter  with  a  solution  of  an  alkali  poly- 
sulphide  and  excess  of  sulphur.  The  latter  method  depends  upon 
the  fact  that  the  black  unstable  modification  is  the  more  soluble  in 
the  polysulphide  solution.  When  the  polysulphide  solution  is  satu- 
rated with  regard  to  the  black  form,  it  is  therefore  supersaturated 
with  regard  to  the  red,  which  separates  out.  Mercuric  sulphide, 
under  the  name  vermilion,  is  used  as  a  pigment. 

Mercuric  sulphide  is  insoluble  in  nitric  or  hydrochloric  acid,  but  is 
dissolved  by  aqua  regia,  soluble  mercuric  chloride  being  formed. 

Compounds  of  Mercury  Salts  with  Ammonia  —  These 
compounds  may  be  regarded  as  being  derived  from  ammonium  salts, 
NH4X,  by  the  substitution  of  part  or  all  of  the  hydrogen  by  Hg2" 
groups  (mercurous  compounds),  or  by  Hg"  atoms  (mercuric  com- 
pounds) ;  or  one  hydrogen  may  be  displaced  from  each  of  two  NH4 
groups.  The  constitution  of  these  compounds  is,  however,  not  well 
understood,  and  other  graphical  formulae  than  those  given  might  be 
suggested. 

Mercurous  Compounds  —  The  black  precipitate  formed  by 
the  addition  of  solution  of  ammonia  to  calomel  is  often  regarded  as 
"mercurous  ammonium  chloride,"  NH2Hg2Cl,  being  derived  from 
ammonium  chloride  by  the  substitution  of  the  Hg2"  group  for  two 
atoms  of  hydrogen  : 


As  a  matter  of  fact,  however,  the  compound  appears  to  be  a  mix- 
ture of  the  corresponding  mercuric  compound  and  finely  divided 
mercury  : 

N  H2Hg2"Cl-»N  H2Hg-Cl  +  Hg. 

Definite  "mercurous  ammonium"  compounds  are  not  known  with 
certainty. 

Mercuric  Compounds.     Mercuric  Ammonium  Chlo- 
ride (Infusible  white  precipitate\  (NH2Hg)Cl,  is  obtained  as  an 


MERCURY  477 

amorphous  powder  when  ammonia  is  added  in  slight  excess  to  a 
solution  of  mercuric  chloride  : 

HgCl2  +  NH3->(NH2Hg)Cl  +  HC1. 

It  decomposes  on  heating  without  previously  fusing. 

A  related  compound,  (NHg2)Cl,  often  mentioned  in  chemical 
text-books,  does  not  appear  to  exist.  The  corresponding  nitrate, 
(NHg2)NO3,  however,  is  obtained  as  a  slightly  yellowish  precipitate 
by  adding  ammonia  to  a  solution  of  mercuric  nitrate  acidified  with 
nitric  acid. 

/NH3C1 

Mercuric  Diammonium  Chloride,  Hg<  ,  or,  better, 

\NH8C1 

Hg(NH3)2Cl2(/'><f.tt£/fc  white  precipitate)^  obtained  by  slowly  adding 
a  solution  of  mercuric  chloride  to  a  boiling  aqueous  solution  of 
ammonium  chloride  and  ammonia  till  the  precipitate  first  formed  no 
longer  dissolves.  On  cooling,  the  compound  separates  in  very  small 
colourless  crystals. 

Oxydimercuric  Ammonium  Iodide,  (OHg2):NH2I,  is 
obtained  by  the  addition  of  ammonia  to  a  solution  of  mercuric  iodide 
in  potassium  iodide,  to  which  excess  of  potassium  hydroxide  has 
been  added  (Nessler's  reagent}.  It  forms  a  brown  powder  with  a 
purple-red  tinge.  It  may  be  regarded  as  a  derivative  of  Milloris 
base,  (OHHg)2  :NH2OH,  which  is  obtained  as  a  yellow  crystalline 
powder  by  the  action  of  an  aqueous  solution  of  ammonia  on  mercuric 
oxide. 

Tests  for  Mercury — Mercury  is  obtained  in  metallic  globules 
by  mixing  any  compound  containing  it  with  sodium  carbonate  and 
heating  in  a  tube.  Further,  copper  becomes  coated  with  mercury 
when  introduced  into  the  solution  of  any  mercury  salt.  Mercurous 
salts  in  solution  give,  with  soluble  chlorides,  a  white  precipitate  of 
calomel,  turned  black  by  ammonia.  The  precipitation  of  the  black 
sulphide  by  means  of  hydrogen  sulphide  is  also  a  useful  test. 

Comparison  of  the  Members  of  the  Zinc  Group, 
and  Summary — The  resemblance  between  the  members  of 
this  sub-group  is  by  no  means  so  close  as  in  other  families  already 
considered.  This  applies  particularly  to  mercury,  which,  as  we« 
have  seen,  differs  in  many  respects  from  the  other  members  of  the 
sub-group. 

Corresponding  with  their  position  in  the  periodic  table,  they  are 
all  divalent  elements.  The  metals  themselves  are  fairly  stable  in 


478     A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 


the  air  at  room  temperature,  but  on  heating  form  the  oxides ;  they 
do  not  decompose  water  at  room  temperature.  They  are  only 
weakly  electro-positive,  and  therefore  the  salts  are  hydrolyzed  in 
solution,  and  there  is  considerable  tendency  to  the  formation  of 
basic  salts.  The  hydroxides  of  beryllium,  zinc  and  cadmium  are  also 
weakly  acidic.  The  more  important  physical  constants  of  the  metals, 
and  their  variation  with  the  atomic  weight,  are  shown  in  the  table. 


Be 

Mg 

Zn 

Cd 

Hg 

Atomic  weight  .         . 

9.1 

24.3 

65.4 

1  12.4 

200.0 

Density     .... 

1.85 

J-75 

7.0 

8.6 

13.6 

Melting-point  .        .         . 

<IOOO° 

800° 

419° 

321 

-39° 

Boiling-point    . 

... 

1100° 

920° 

780 

357° 

Atomic  volume 

4.9 

13.9 

9-3 

13-1 

14.7 

The  chemical  properties  are  of  course  determined  mainly  by  the 
weakly  electro-positive  character  of  the  metals.  Omitting  beryllium) 
the  behaviour  of  which  is  uncertain,  the  electro-positive  character 
diminishes  in  the  order  :  magnesium,  zinc,  cadmium,  mercury  (p.  434), 
but  the  basic  character  does  not  diminish  regularly  with  increasing 
atomic  weight  (cf.  p.  388).  Magnesium  and  zinc  resemble  each  other 
very  closely,  but  differ  in  the  important  point  that  zinc  and  its 
hydroxide  are  soluble,  magnesium  and  its  hydroxide  insoluble,  in 
alkalis.  Further,  whereas  magnesium  sulphate  in  aqueous  solution  is 
practically  neutral  to  litmus,  zinc  sulphate  solution  is  acid  owing  to 
hydrolysis.  Magnesium  is,  in  fact,  a  considerably  stronger  base  (as 
hydroxide)  than  zinc.  In  its  behaviour  towards  -alkalis,  beryllium 
resembles  zinc  and  also  aluminium,  and  is  therefore  less  basic  than 
magnesium.  Cadmium  closely  resembles  zinc,  but  its  hydroxide  is 
insoluble  in  alkali  and  is  therefore  probably  a  stronger  base  than  zinc 
hydroxide. 

Mercury  resembles  the  other  elements  of  the  sub-group  fairly 
closely  in  the  mercuric  compounds,  but  differs  in  forming  mercuric 
compounds  containing  the  divalent  Hg2"  group.  The  mercurous 
compounds  are  in  some  respects  analogous  to  the  cuprous  compounds 
'and  the  mercuric  to  the  cupric  compounds.  The  slight  degree  of 
ionisation  of  the  halides  of  cadmium  and  of  divalent  mercury  is 
noteworthy. 


CHAPTER   XXX 
ELEMENTS  OF  ALUMINIUM  GROUP  (GROUP   III) 

Sub-group  A  Sub-group  B 

Scandium  (Sc)  .  .  .  .44.1  Boron  (B)  .  .  .  .  u.o 
Yttrium  (Yt)  ....  89.0  Aluminium  (Al)  .  .  .27.1 
Lanthanum  (La)  .  .  .  139.0  Gallium  (Ga)  .  .  .  .69.9 
Ytterbium  (Yb) .  .  .  .  172.0  Indium  (In)  ....  114.8 

Thallium  (Tl)    ....     204.0 

WITH  the  exception  of  boron,  aluminium,  and  thallium,  the  mem- 
bers of  the  third  group  occur  in  extremely  small  proportion  in 
nature.  Corresponding  with  their  position  in  the  periodic  table,  all  are 
tervalent  elements.  The  oxides  are  therefore  of  the  type  M2O3.  That 
of  boron  (p.  375)  is  acidic,  the  others  are  weakly  basic,  but  aluminium 
oxide  has  also  weak  acidic  properties.  Besides  the  tervalent  (thallic) 
compounds,  thallium  forms  a  series  of  compounds  (the  thallous  com- 
pounds) in  which  it  functions  as  a  univalent  element. 

Boron  has  already  been  considered  in  connexion  with  the  non- 
metals  (p.  372).  From  considerations  of  space,  the  rare  elements  are 
dealt  with  very  briefly ;  for  details,  larger  works  on  the  subject  must 
be  consulted. 

ALUMINIUM 

Symbol,  Al.     Atomic  weight=27.i. 

Occurrence — Aluminium  does  not  occur  free  in  nature,  but  in 
the  combined  condition  is  an  extremely  abundant  element  (p.  19). 
The  oxide,  A12O3,  occurs  in  crystalline  form  in  corundum  or  emery, 
an  extremely  hard  substance  used  for  polishing,  and  also  in  ruby  and 
sapphire;  all  these  substances  are  coloured  by  traces  of  impurities. 
As  hydrated  oxide,  A12O3,2H2O,  it  occurs,  along  with  iron  oxide  and 
silicon,  in  bauxite  (found  in  France  and  the  United  States),  which  is 
an  important  commercial  source  of  the  metal.  As  fluoride,  associated 
with  sodium  fluoride,  it  occurs  in  large  amount  as  cryolite,  A1F3,  3NaF, 
in  Greenland.  As  silicate,  it  is  an  important  constituent  of  felspar, 
mica,  and  the  other  compounds  of  which  the  silicate  rocks  (granite, 

479 


480     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

etc.)  are  built  up.  The  different  clays,  which  result  from  the  dis- 
integration of  rocks,  are  largely  composed  of  aluminium  silicate 
(P-  37°)-  Although  a  regular  constituent  of  soils,  it  is  not  taken  up 
to  any  extent  by  plants. 

Preparation  of  Metal  —  Formerly,  aluminium  was  prepared 
by  the  action  of  metallic  sodium  on  the  double  chloride  of  aluminium 
and  sodium  : 

AlCl3,NaCl  +  3Na-»Al  +  4NaCl. 

It  is  now  prepared  exclusively  by  the  electrolysis  of  aluminium  oxide 
in  a  bath  of  fused  cryolite. 

The  aluminium  oxide  used  for  this  purpose  must  be  practically  free 
from  iron  and  silica.  When  bauxite  is  used  as  the  source  of  the 
oxide,  it  is  first  fused  with  sodium  carbonate  : 

A12O3  +  3Na2CO3-»2Na3AlO3  +  3CO2, 

the  sodium  aluminate,  Na3AlO3,  is  extracted  with  water  (the  ferric  oxide 
remaining  undissolved),  and  carbon  dioxide  passed  into  the  solution, 
when  aluminium  hydroxite  is  precipitated  : 


The  hydroxide  is  then  washed,  dried,  and  ignited. 

A  simpler  method  of  purification  is  to  dissolve  bauxite  up  to 
saturation  in  a  solution  of  sodium  hydroxide  and  then  stir  crystal- 
line aluminium  oxide  into  the  solution.  As  the  latter  is  greatly  super- 
saturated with  reference  to  the  crystalline  hydroxide,  the  greater 
portion  of  the  dissolved  hydroxide  separates  out. 

The  electrolysis  is  carried  out  in  iron  vessels  lined  in  the  interior 
with  carbon  ;  the  vessel  itself  forms  the  cathode,  and  carbon  rods 
are  used  as  anode.  "The  electrolyte,  usually  a  mixture  of  cryolite  and 
aluminium  fluoride,  is  kept  fused  by  the  current,  and  aluminium 
oxide  is  added  from  time  to  time  as  electrolysis  proceeds. 

Properties  —  Aluminium  is  a  white  metal  of  density  2.7  ;  it  melts 
at  657°,  and  is  not  volatile  at  a  white  heat.  It  is  very  ductile  and 
malleable  ;  it  can  readily  be  drawn  into  wire  and  beaten  into  sheets. 
It  is  very  tenacious  under  ordinary  conditions,  but  at  high  tempera- 
tures becomes  brittle.  It  is  an  excellent  conductor  of  heat  and 
electricity. 

At  room  temperature,  aluminium  becomes  coated  with  a  thin  film 
of  oxide,  which  protects  it  against  further  oxidation.  At  high  tem- 
peratures, however,  it  burns  in  air  to  the  oxide  with  a  brilliant  flame 


ALUMINIUM  481 

and  the  evolution  of  a  large  amount  of  heat.  Pure  aluminium  is  only 
very  slightly  acted  on  by  water  or  steam,  but  the  commercial  metal  is 
attacked  by  these  reagents,  doubtless  owing  to  the  formation  of  local 
currents  set  up  between  the  metal  and  the  impurities.  The  fact  that 
aluminium  is  a  highly  electro-positive  metal  is  of  importance  in  this 
connexion  (cf.  p.  434). 

Dilute  sulphuric. acid  has  very  little  action  on  aluminium,  but  the 
metal  is  dissolved  by  the  concentrated  acid  with  formation  of 
aluminium  sulphate  and  evolution  of  sulphur  dioxide.  Hyjdrpchlpnc, 
acid  readily  dissolves  aluminium,  but  nitric  acid,  whether  dilute  or 
concentrated,  is  practically  without  action  at  room  temperature, 
although  at  the  boiling-point  a  vigorous  reaction  takes  place.  Phos: 
phoric  acid^  both  in  dilute  and  concentrated  solution,  readily  dis- 
solves aluminium.  Organic  acids  (e.g.  acetic  acid)  are  almost 
without  action  on  aluminium  at  room  temperature,  but  have  a  slight 
solvent  action  in  the  presence  of  sodium  chloride.  Aluminium  is 
dissolved  by  alkali  hydroxides  with  formation  of  aluminates  and 
evolution  of  hydrogen  : 


Uses  of  Aluminium  Alloys— On  account  of  its  lightness, 
relative  cheapness  and  high  electrical  conductivity,  aluminium  is 
being  used  instead  of  copper  for  the  conveyance  of  electricity.  It 
is  also  used  in  making  cooking  utensils,  pans,  trays,  etc.,  but  the  high 
expectations  formerly  held  as  to  its  usefulness  have  scarcely  been 
realized  in  practice,  owing  to  the  fact,  already  mentioned,  that  it  is 
attacked  by  alkalis,  including  soap  solution,  and  also,  unless  very 
pure,  by  water  and  by  acetic  acid. 

The  finely  divided  metal,  mixed  with  oil,  is  used  as  a  paint. 
Aluminium  coated  with  mercury  (aluminium  amalgam),  prepared  by 
dipping  sheet  aluminium  into  a  solution  of  mercuric  chloride,  decom- 
poses water  at  room  temperature,  with  liberation  of  hydrogen.  The 
greater  activity  of  the  amalgam,  as  compared  with  aluminium,  is  partly 
at  least  due  to  the  fact  that  the  mercury  prevents  the  formation  of  a 
protective  coating  of  oxide  on  the  surface  of  the  metal. 

Certain  other  alloys  of  aluminium  are  in  use.  Aluminium  bronze 
(copper  with  5  to  12  per  cent,  of  aluminium)  is  golden-yellow  in 
colour,  very  tough,  takes  an  excellent  polish,  and  gives  good  castings. 
Magnalium  (aluminium  with  10  to  25  per  cent,  of  magnesium)  is  also 
used  ;  it  is  harder  than  aluminium  and  takes  a  good  polish. 

The  great  affinity  between  aluminium  and  oxygen,  already  referred 
31 


482     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

to,  is  illustrated  by  the  equation  representing  the  heat  of  the  forma- 
tion of  the  oxide  : 


x  380,000  cal. 

The  use  of  aluminium  for  obtaining  other  metals  from  their  oxides, 
according  to  the  so-called  "  thermite  "  process  (Goldschmidt),  is  based 
on  this  fact.  The  oxide  in  'question  (e.g.  chromium,  iron  or  man- 
ganese oxide)  is  mixed  with  finely  divided  aluminium,  and  the  reaction 
initiated  by  a  fuse  of  magnesium  ribbon.  It  then  proceeds  with  the 
evolution  of  much  heat,  the  temperature,  under  favourable  conditions, 
rising  to  3000°  ;  the  final  products  are  the  metal  and  aluminium  oxide. 
The  same  process  is  used  for  welding  pieces  of  metal,  e.g.  steel  rails. 
The  junction  is  surrounded  by  a  mixture  of  ferric  oxide  and  aluminium, 
which  is  ignited  in  the  usual  way,  and  under  the  influence  of  the  high 
temperature  the  two  pieces  of  metal  are  fused  and  thoroughly  welded 
together. 

Aluminium  Oxide  (Alumina),  A12O3,  occurs  naturally  in  anhy- 
drous rhombohedral  crystals  as  corundum,  which  is  colourless  ;  as 
ruby,  coloured  by  traces  of  chromium  salts  ;  as  sapphire,  coloured  blue 
by  cobalt  salts,  and  as  amethyst.  Emery  is  aluminium  oxide  contain- 
ing ferric  oxide  ;  owing  to  its  great  hardness  it  is  used  for  polishing. 
The  hydrated  oxide  occurs  naturally  as  bauxite,  A12O3,3H2O,  and  as 
diaspore,  A12O3,H2O. 

The  oxide  is  obtained  as  a  white  amorphous  powder  by  heating 
aluminium  hydroxide,  A1(OH)3  ;  in  this  form  it  dissolves  readily  in 
acids  and  in  alkalis.  When  the  oxide  is  strongly  ignited  it  becomes 
practically  insoluble  in  acids  ;  this  is  probably  due,  in  part  at  least, 
to  conversion  into  the  insoluble  crystalline  form.  The  insoluble  oxide 
is  brought  into  solution  by  fusing  with  alkalis  (p.  480). 

Aluminium  Hydroxide,  A1(OH)3,  is  obtained  as  a  colourless 
gelatinous  precipitate  (hydrogel,  p.  371)  by  adding  an  alkali  hy- 
droxide (not  in  excess)  or  ammonium  hydroxide  to  the  solution  of 
an  aluminium  salt  : 

A1C13+3NH4OH->A1(OH)3  +  3NH4C1. 

The  precipitate  thus  obtained  contains  more  or  less  adsorbed  water 
(p.  329).  On  heating,  the  water  is  driven  off  and  finally,  on  strong 
heating,  the  anhydrous  oxide  is  obtained.  The  hydrates,  A12O3,3H2O 
(  =  A1(OH)3),  A12O3,2H2O,  and  A12O3,H2O,  are  presumably  successive 
stages  in  the  dehydration  of  the  gelatinous  hydroxide. 
Aluminium  hydroxide  is  at  the  same  time  a  weak  base  and  a  weak 


ALUMINIUM  483 

acid.  The  basic  character  is  indicated  by  the  existence  of  salts  such 
as  aluminium  chloride  and  aluminium  sulphate  ;  and  the  fact  that  the 
hydroxide  is  a  weak  base  is  shown  by  the  partial  hydrolysis  of  these 
salts  in  solution,  and  by  the  fact  that  salts  with  weak  acids  (e.g. 
carbonate,  sulphide)  cannot  be  obtained  in  the  presence  of  water. 
The  weakly  acidic  character  of  the  hydroxide  is  indicated  by  the 
formation  of  salts,  the  aluminates,  with  strong,  but  not  with  weak 
bases.  This  is  illustrated  by  the  solubility  of  the  hydroxide  in  excess 
of  potassium  hydroxide  : 

A1(OH)3+3KOH->A1(OK)34-3H20, 

and  its  insolubility  in  ammonium  hydroxide  solution.  The  acid 
H3AlO3(or  HA1O2)  is  so  weak  that  potassium  aluminate  ishydrolyzed 
to  a  considerable  extent  in  solution,  and  is  decomposed  even  by 
carbonic  acid,  with  precipitation  of  aluminium  hydroxide  (p.  480). 
The  compound  HA1O2,  derived  from  the  normal  acid  by  abstraction 
of  a  molecule  of  water,  may  be  termed  meta-aluminic  acid  (cf.  silicic 
acids,  p.  369).  The  minerals  spinelle,  Mg(AlO2)2,  and  chrysoberyl, 
Be(AlO2)2,  may  be  regarded  as  being  derived  from  the  latter  acid. 

Aluminium  hydroxide  may  be  obtained  in  colloidal  solution 
(hydrosol)  by  dissolving  the  hydrogel  in  a  solution  of  aluminium 
chloride  and  then  removing  the  latter  by  dialysis  (p.  369). 

When  aluminium  hydroxide  is  precipitated  in  a  solution  containing 
a  colouring  matter,  the  latter  is  carried  down  with  the  precipitate, 
leaving  the  liquid  almost  colourless.  In  such  a  case  the  hydroxide 
is  said  to  have  adsorbed  the  colouring  matter.  The  precipitates  are 
termed  lakes. 

Advantage  is  taken  of  this  fact  in  causing  certain  dyes  to  adhere 
firmly  to  cloth.  The  aluminium  hydroxide  is  first  precipitated  in  the 
fibres  of  the  cloth,  which  is  then  immersed  in  the  dye.  The  latter  is 
removed  from  the  solution  by  the  aluminium  hydroxide  and  "  fixed " 
on  the  cloth.  Substances  used  for  the  purpose  of  fixing  dyes  are 
known  as  mordants.  Some  dyes  unite  directly  with  the  fibre,  and  in 
such  cases  mordants  are  unnecessary. 

When  a  substance  is  at  the  same  time  a  weak  acid  and  a  weak  base, 
both  properties  are  necessarily  weak.  That  this  must  be  so  is  clear 
when  it  is  remembered  that  the  acidic  and  basic  characters  depend 
upon  the  presence  of  H'  and  OH'  ions  respectively,  and  that  both 
cannot  be  present  in  considerable  proportion  in  the  same  solution 
owing  to  their  tendency  to  unite  to  form  water. 


484     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Aluminium  Chloride,  A1C13,  is  obtained  in  the  anhydrous 
form  by  heating  aluminium  foil  in  dry  chlorine  or  hydrogen  chloride, 
or  by  strongly  heating  a  mixture  of  aluminium  oxide  and  charcoal  in 
a  current  of  chlorine  ;  the  chloride  is  collected  in  a  dry  receiver.  It 
is  obtained  in  aqueous  solution  by  dissolving  aluminium  oxide  in 
hydrochloric  acid;  on  evaporating  the  solution  the  hexahydrate, 
A1C13,6H2O,  separates  in  colourless  crystals. 

The  anhydrous  chloride  occurs  in  colourless  crystals,  which  fume  in 
moist  air;  it  vaporizes,  without  melting,  at  183°.  At  low  tempera- 
tures the  vapour  density  corresponds  with  the  formula  A12C16,  but 
above  835°  it  is  present  entirely  as  A1C13  molecules.  It  is  readily 
soluble  in  water,  and  undergoes  considerable  hydrolysis  ;  when  the 
aqueous  solution  is  evaporated  to  dryness  the  residue  consists  almost 
entirely  of  the  hydroxide. 

Aluminium  chloride  is  largely  used  as  a  catalytic  agent  in  organic 
chemistry. 

Aluminium  Sulphate,  A12(SO4)3,  is  obtained  on  the  large 
scale  by  dissolving  bauxite  in  sulphuric  acid  (the  product  contains 
iron  as  impurity),  and  also  by  heating  finely-divided  clay  (aluminium 
silicate)  with  concentrated  sulphuric  acid.  The  liquid  is  separated 
from  the  insoluble  silica  and  other  impurities  and  concentrated,  when 
the  salt  separates  as  Al2(SO4)3,i8H2O  in  monoclinic  crystals.  The 
pure  salt  is  obtained  by  dissolving  the  pure  oxide  or  hydroxide  in 
sulphuric  acid. 

The  salt  with  i8H2O  becomes  anhydrous  on  heating,  and  at  a 
higher  temperature  is  completely  decomposed  into  aluminium  oxide 
and  sulphur  trioxide.  It. is  readily  soluble  in  water  and  the  aqueous 
solution  is  acid  owing  to  hydrolysis : 

A12(S04)3  +  2H20$A12(OH)2(S04)2  +  H2S04. 

Aluminium  sulphate  is  used  as  a  mordant  (p.  483). 

Alums — Aluminium  sulphate  forms  with  the  alkali  sulphates  a 
series  of  double  salts  known  as  alums,  of  which  potassium  alum, 
A12(SO4)3,K2SO4,24H2O  (or  A1K(SO4)2,I2H2O),  is  a  type.  The  place 
of  the  potassium  may  be  talcen  by  sodium,  rubidium,  caesium, 
ammonium,  silver,  or  thallium,  and  further,  the  isomorphous  salts, 
which  contain  trivalent  chromium  or  iron  instead  of  aluminium,  are 
also  termed  alums.  The  alums  all  form  octahedral  crystals  with 
24H<>O,  and  are  isomorphous.  The  general  formula  for  any  alum  is 
from  the  above : 

M2III(S04)3,M2IS04,24H20  or  MIIIMI(SO)2,i2H2O. 


ALUMINIUM  485 

When  the  alums  are  strongly  heated,  they  lose  water  and  sulphur 
trioxide,  and  finally  yield  a  mixture  of  aluminium  oxide  and  alkali 
sulphate.  When  ammonium  alum  is  heated  the  final  product  is 
aluminium  oxide. 

The  alums  are  fairly  soluble  in  water,  and  in  dilute  solution  are 
decomposed  into  the  component  salts.  The  solutions  have  an  acid 
reaction,  owing  to  partial  hydrolysis.  They  are  used  as  mordants. 

Potassium  alum,  the  most  important  of  the  alums,  was  known  to  the 
ancient  Romans.  It  is  prepared  from  the  naturally  occurring  basic 
sulphate,  alunite,  A12(SO4)3,K2SO4,4A1(OH)3,2H2O,  by  roasting,  ex- 
tracting with  water,  whereby  alumina  remains  undissolved,  and  then 
concentrating  the  solution.  It  may  also  be  prepared  by  the  action 
of  sulphuric  acid  on  bauxite  or  on  clay  ;  to  the  aluminium  sulphate 
thus  obtained'the  requisite  amount  of  potassium  sulphate  is  added. 

Potassium  alum  loses  all  its  water  on  heating  to  100°  in  a  current 
of  air.  The  product  dehydrated  at  a  fairly  high  temperature  is 
termed  burnt  alum;  it  is  usually  slightly  basic  owing  to  the  loss  of 
sulphur  trioxide. 

The  solubility  of  potassium  alum  in  water  increases  rapidly  with  rise 
of  temperature  ;  100  grams  of  water  take  up  at  o°  3.9  grams,  at  20° 
15.1  grams,  at  50°  44.1  grams,  at  100°  357  grams  of  the  hydrated  salt. 

Ammonium  alum  is  obtained  in  well-formed  crystals,  and  is  a 
familiar  salt.  Sodium  alum,  on  the  other  hand,  crystallizes  with 
difficulty,  and  is  very  little  used. 

Another  class  of  double  sulphates,  sometimes  termed  pseudo-alums^ 
are  known;  they  differ  from  the  alums  in  containing  one  atom  of 
a  divalent  metal,  e.g.  manganese,  copper,  zinc,  iron,  or  magnesium, 
instead  of  two  atoms  of  a  univalent  element.  The  aluminium-iron 
pseudo-alum  has  the  formula  Al2(SO4)3,FeSO4,24H2O.  They  are  not 
isomorphous  with  the  alums. 

Aluminium  Silicates — Reference  has  already  been  made  to 
the  fact  that  double  silicates  containing  aluminium  are  important 
constituents  of  rocks  (granite,  etc.).  When  these  rocks  are  dis- 
integrated by  moisture  and  carbon  dioxide,  aluminium  silicate,  being 
comparatively  insoluble,  remains,  and  the  other  constituents  are 
dissolved  and  conveyed  into  the  soil.  The  purest  form  of  aluminium 
silicate  is  kaolin,  Al2O3,2SiO2,2H2O,  which  is  colourless  and  is  used 
in  making  porcelain.  Ordinary  clay  also  consists  largely  of  aluminium 
silicate,  but,  unlike  kaolin,  it  is  a  sedimentary  deposit,  and  contains 
ferric  oxide  (to  which  the  brown  colour  is  due),  sand,  and  calcium 
carbonate  as  impurities. 


486    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

When  mixed  with  water,  clays  form  a  plastic  mass,  which  can  be 
moulded  into  any  desired  shape  and  sets  hard  on  heating  strongly 
(burning  or  "  firing").  On  this  depends  their  use  in  the  manufacture 
of  bricks,  earthenware  and  porcelain.  Bricks  are  made  from  impure 
clay  containing  ferric  oxide  and  calcium  carbonate.  Earthenware 
articles  are  prepared  in  the  same  way  from  rather  purer  clays  ;  the 
surface  is  usually  covered  by  a  glaze  of  fused  silicate  obtained  by 
throwing  sodium  chloride  into  the  vessel  while  in  process  of  firing. 
Porcelain  is  made  from  pure  clay  (kaolin),  and  the  pores  are  com- 
pletely filled  with  fused  silicate,  obtained  by  adding  felspar  and 
quartz  before  burning. 

The  use  of  clay  for  making  bricks,  etc.,  was  known  to  the  ancient 
Babylonians  and  Egyptians  at  least  3000  years  B.C.,  and  our  know- 
ledge of  the  early  history  of  the  Babylonians  depends  upon  the  fact 
that  their  records  (letters,  lists  of  laws,  etc.)  were  impressed  upon 
moist  clay,  which  was  afterwards  fired. 

Ultramarine  is  a  blue  pigment  obtained  by  heating  together 
kaolin,  sodium  carbonate,  sodium  sulphate,  sulphur  and  carbon.  It 
occurs  in  nature  as  lapis  lazuli.  It  is  usually  regarded  as  a  double 
silicate  of  aluminium  and  sodium  associated  with  sodium  poly- 
sulphides,  but  its  constitution  is  not  yet  fully  understood.  It  is  decom- 
posed by  acids  with  evolution  of  hydrogen  sulphide  and  disappearance 
of  the  blue  colour.  It  is  largely  used  to  neutralize  the  yellow  tint  in 
sugar  and  in  linen  and  cotton  goods.  It  is  a  constituent  of  laundry 
blue. 

Aluminium  Sulphide,  A12S3,  is  obtained  as  a  black  amor- 
phous powder  by  adding  sulphur  to  melted  aluminium.  It  is  com- 
pletely decomposed  by  water  with  formation  of  aluminium  hydroxide 
and  hydrogen  sulphide  : 


and  burns  in  air  to  the  oxide  and  sulphur  dioxide.  As  both  base  and 
acid  are  very  weak,  it  cannot  be  obtained  by  interaction  in  aqueous 
solution  owing  to  hydrolysis. 

Tests  for  Aluminium  —  The  formation  of  a  white  gelatinous 
precipitate  when  an  alkali  is  added  to  the  solution  of  an  aluminium 
salt,  the  precipitate  being  soluble  in  excess  of  potassium  or  sodium 
hydroxide,  but  insoluble  in  ammonia.  Further,  when  ammonium 
chioride  is  added  to  a  solution  of  potassium  aluminate,  aluminium 
hydroxide  is  precipitated  ;  this  test  depends  upon  the  formation  and 
subsequent  hydrolysis  of  ammonium  aluminate.  When  aluminium 


THALLIUM  487 

oxide  is  moistened  with  a  few  drops  of  cobalt  nitrate  solution  and 
strongly  heated  on  charcoal  in  the  blowpipe  flame,  a  blue  mass  is 
obtained. 

GALLIUM 
Symbol,  Ga.     Atomic  weight=69.g 

Gallium  was  discovered  in  1875  by  Lecoq  de  Boisbaudran  in  zinc  blende  by 
means  of  the  spectroscope.  Its  existence  and  properties  had  been  foretold  some 
years  before  by  Mendel6eff  on  the  basis  of  the  Periodic  System  (p.  369).  Gallium 
is  a  very  rare  element. 

The  metal  itself  is  white  and  melts  at  30°  ;  its  density  is  5.9.  In  most  of  its 
properties  it  strongly  resembles  aluminium.  Thus  it  has  no  action  on  water,  it 
becomes  coated  with  a  film  of  oxide  on  heating  in  the  air,  it  dissolves  readily  in 
hydrochloric  acid  and  in  alkali  hydroxides,  but  is  insoluble  in  dilute  nitric  acid  at 
room  temperature. 

In  its  most  important  compounds  gallium  is  trivalent.  The  hydroxide, 
Ga(OH)3,  is  soluble  in  alkalis,  the  chloride,  GaCl3,  fumes  in  the  air,  and  is 
hydrolyzed  in  aqueous  solution,  and  the  sulphate  forms  a  well-defined  alum, 
Ga2(SO4)3,(NH4)2SO4,24H2O,  with  ammonium  sulphate. 

INDIUM 
Symbol,  In.     Atomic  weight= 114.8. 

Indium  is  also  a  very  rare  element,  and  shows  many  analogies  with  aluminium. 
It  was  discovered  in  1863  by  Reich  and  Richter  in  a  zinc  ore  from  Freiberg  by 
means  of  its  spectrum,  which  is  characterized  by  the  presence  of  a  bright  blue 
line.  The  metal  itself  is  colourless  and  very  soft;  it  melts  about  155°.  It  is 
stable  in  the  air  at  room  temperature,  but  on  heating  burns  with  a  blue  flame  to 
the  oxide,  In2O3.  It  slowly  decomposes  water  at  room  temperature  with  forma- 
tion of  the  hydroxide. 

In  its  most  important  series  ot  compounds  indium  is  trivalent.  The  hy- 
droxide, In(OH)3,  is  soluble  in  excess  of  potassium  or  sodium  hydroxide,  so  that 
it  has  weakly  acidic  as  well  as  basic  properties.  The  chloride,  InCl3,  unlike 
those  of  gallium  and  aluminium,  can  be  obtained  without  decomposition  by 
evaporation  of  its  aqueous  solution  at  100°.  The  sulphate,  In2(SO4)3,  forms  an 
alum  with  ammonium  sulphate. 

Besides  acting  as  a  trivalent  element,  indium  forms  two  series  of  compounds  in 
which  it  functions  as  a  univalent  and  as  a  divalent  element  respectively.  The 
halogen  compounds  of  these  series,  e.g.  InCl  and  InCl2,  are  best  known. 

THALLIUM 

Symbol,  Tl.     Atomic  weight =204. 6. 

Thallium  was  discovered  by  Crookes  in  1861  in  the  mud  of  a 
sulphuric  acid  chamber  at  Tilkerorde  in  the  Hartz  mountains  by 
means  of  the  spectroscope.  The  spectrum  is  characterized  by  the 


488    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

presence  of  a  bright  green  line,  hence  the  name  (from  6a\\6s,  a  green 
twig).  It  is  also  found  in  the  flue  dust  of  the  chambers.  The  only 
mineral  which  contains  a  considerable  proportion  of  thallium  is 
crookesite,  a  selenide  of  copper  which  contains  16  to  18  per  cent,  of 
thallium  and  3  to  5  per  cent,  of  silver. 

Preparation  and  Properties  of  Metal — Thallium  is 
best  obtained  from  the  dust  in  the  flues  of  the  sulphuric  acid 
chamber  by  boiling  with  water,  pouring  off  the  solution,  and  adding  to 
it  sodium  chloride,  which  precipitates  the  thallium  as  T1C1.  The 
chloride  is  further  purified,  heated  with  sulphuric  acid  to  obtain  the 
sulphate,  and  the  aqueous  solution  of  the  latter  electrolyzed  with 
a  platinum  anode  and  a  cathode  of  copper  foil. 

Thallium  is  a  soft,  heavy  metal  with  a  bluish  tinge,  like  lead. 
Its  density  is  n.8;  it  melts  at  285°  and  boils  at  1650°.  At  room 
temperature  it  rapidly  becomes  coated  with  a  mixture  of  oxides. 
It  is  stable  in  air-free  water  at  ordinary  temperature,  but  decomposes 
water  at  a  red  heat.  The  metal  is  slowly  acted  upon  by  water 
containing  dissolved  oxygen,  with  formation  of  thallous  hydroxide, 
Tl(OH),  which  is  soluble  in  water.  When  heated  in  air,  it  burns 
to  the  trioxide  with  a  green  flame.  It  is  readily  soluble  in  dilute 
nitric  or  sulphuric  acids,  but  is  only  acted  on  slowly  by  hydrochloric 
acid,  owing  to  the  insolubility  of  the  chloride,  T1C1. 

Compounds  of  Thallium — Thallium  forms  two  well-defined 
series  of  compounds,  thallous  compounds,  in  which  it  is  univalent, 
and  thallic  compounds,  in  which  it  is  trivalent.  The  thallous  salts 
resemble  those  of  the  alkalis  inasmuch  as  the  hydroxide  and  car- 
bonate are  soluble  in  water,  whilst  the  relatively  slight  solubility 
of  the  thallous  halides  recalls  the  silver  halides.  In  its  trivalent 
compounds,  thallium  resembles  aluminium  and  gold. 

Thallous  Compounds—  Thallous  oxide,  T12O,  is  obtained 
as  a  black,  hygroscopic  powder  by  heating  thallous  hydroxide  to 
100°.  Thallous  hydroxide,  T1OH,  is  obtained  by  the  simul- 
taneous action  of  water  and  oxygen  on  the  metal  or  by  double 
decomposition  between  barium  hydroxide  and  thallous  sulphate 
in  solution.  It  is  soluble  in  water  and  is  a  fairly  strong  base. 
Thallous  chloride,  T1C1,  is  obtained  as  a  white  curdy  precipi- 
tate by  adding  a  soluble  chloride  to  the  solution  of  a  thallous 
salt.  At  20°  TOO  grams  •  of  water  take  up  0.34  grams,  at  100° 
2.4  grams  of  the  salt.  Thallous  iodide,  Til,  is  obtained  by 
double  decomposition  as  a  k  yellow  precipitate.  At  20°  100  grams 
of  water  dissolve  0.063  grams  of  the  salt.  Thallous  carbonate, 


THALLIUM  489 

T12CO3,  is  obtained  by  saturating  a  solution  of  thallous  hydroxide 
with  carbon  dioxide.  It  is  fairly  soluble  in  water  (about  5  grams 
in  100  grams  of  water  at  18°),  and  separates  in  long,  prismatic 
needles  on  evaporating  the  solution. 

Thallic  Compounds—  Thallic  oxide,  T12O3,  is  obtained  as 
a  dark  brown  powder  by  burning  thallium  in  the  air,  or  by  heating 
the  hydroxide,  T1(OH)3,  or  hydrated  oxide,  T12O3,H2O.  It  is  in- 
soluble in  water,  is  reduced  by  heating  with  hydrochloric  acid 
with  evolution  of  chlorine,  and  is  also  reduced  by  boiling  with 
sulphuric  acid,  oxygen  being  given  off  and  thallous  sulphate  formed. 
At  a  red  heat  thallic  oxide  is  completely  decomposed  into  thallous 
oxide  and  oxygen.  Thallic  hydroxide,  T1(OH)3  (  =  T12O3,3H2O),  is 
said  to  be  formed  as  a  brown  precipitate  by  the  hydrolytic  decom- 
position of  thallic  salts,  but  the  compound  thus  obtained  may  be 
T12O3,H2O.  It  is  known  that  the  voluminous  reddish-brown  pre- 
cipitate obtained  on  adding  ammonia  to  a  solution  of  a  thallic 
salt  has  the  composition  T12O3,H2O.  Thallic  hydroxide  is  a  much 
weaker  base  than  thallous  hydroxide.  Thallic  chloride,  T1C13,  is 
obtained  by  passing  chlorine  into  thallous  chloride  suspended  in 
water  till  a  clear  solution  is  obtained  ;  on  evaporation  the  tetra- 
hydrate,  T1C13,4H2O,  separates  in  colourless  needles.  By  dehydrating 
in  a  vacuum  over  sulphuric  acid,  the  anhydrous  salt  is  finally  obtained. 
The  aqueous  solution  of  the  trichloride  has  an  acid  reaction  owing 
to  hydrolysis,  and  when  it  is  considerably  diluted  the  oxyhydrate, 
T12O3,H2O,  is  precipitated.  Thallic  iodide,  T1I3,  is  obtained  in 
black,  rhombic  crystals  by  the  action  of  iodine  in  alcoholic  solution 
on  thallous  iodide.  It  is  isomorphous  with  the  tri-iodides  of  rubidium 
and  caesium  (p.  418).  It  very  readily  splits  up  into  thallous  iodide  and 
iodine  : 


Thallic  sulphate,  T12(SO4)3,  has  been  obtained  in  the  anhydrous 
form  by  heating  thallium  hydrogen  sulphate,  T12(SO4)3,H2SO4,8H2O, 
at  220°.  The  latter  salt  is  obtained  by  dissolving  thallic  oxide 
in  dilute  sulphuric  acid  and  evaporating  the  solution.  Thallic 
nitrate,  T1(NO3)3,  is  obtained  in  colourless  lustrous  crystals  with 
3H2O  by  dissolving  thallic  oxide  in  concentrated  nitric  acid  and 
evaporating  the  solution.  It  is  hydrolyzed  by  water  with  precipita- 
tion of  thallic  oxide  monohydrate. 

A  number  of  double  thallous-thallic  salts,  of  the  type 
are  known. 


490    A  TEXT-BOOK   OF  INORGANIC   CHEMISTRY 

Thallium  salts  are  readily  recognized  by  their  characteristic 
spectrum. 

THE  SCANDIUM  SUB-GROUP 

As  already  mentioned,  this  sub-group  consists  of  four  rare  elements,  scandium, 
yttrium,  lanthanum,  ytterbium.  They  belong  to  the  so-called  rare  earths,  a 
group  of  elements  which  are  so  similar  in  behaviour  that  very  great  difficulty  has 
been  experienced  in  separating  them.  Fractional  precipitation  and  fractional 
crystallization  have  been  largely  used  for  this  purpose,  and  advantage  has  also 
been  taken  of  the  fact  that  the  nitrates  decompose  at  different  temperatures  on 
heating.  Most  of  them  have  characteristic  spectra. 

Scandium  (Sc= 44.  i )  was  discovered  in  1879  by  Nilson  and  by  Cleve  in  certain 
rare  minerals — gadolinite  and  euxenite — found  in  Scandinavia.  Its  properties 
were  found  to  correspond  closely  with  those  of  the  till  then  unknown  element 
ekaboron,  as  foretold  by  Mendeleeff.  It  forms  only  one  oxide,  Sc2O3,  which 
occurs  as  a  bulky  white  powder;  and  a  corresponding  series  of  salts.  The 
hydroxide,  Sc(OH)3,  occurs  as  a  white  gelatinous  mass,  and  is  a  very  weak  base. 

Yttrium  was  discovered  by  Mosander  (1843)  and,  like  scandium,  occurs  in  the 
gadolinite  earths.  It  always  functions  as  a  trivalent  element.  The  oxide, 
Y2O3,  is  colourless,  and  dissolves  in  acids  to  form  corresponding  salts.  Like 
the  other  rare  earths,  it  gives  a  characteristic  spectrum,  consisting  of  numerous 
lines. 

Lanthanum,  (La=  139.0)  was  discovered  in  1839  in  the  cerite  earths  by 
Mosander.  The  metal  is  obtained  by  electrolysis  of  the  fused  chloride.  It 
is  malleable,  becomes  coated  with  oxide  in  dry  air  and  burns  vigorously  to 
a  mixture  of  oxide  and  nitride  on  heating.  It  decomposes  water  at  room 
temperature. 

Lanthanum  is  invariably  trivalent  in  its  compounds.  The  oxide,  La2O3,  is 
colourless,  and  is  the  strongest  base  among  the  rare  earths.  The  salts  are  also 
colourless,  and  are  only  very  slightly  hydrolyzed  in  aqueous  solution. 

Ytterbium  (Yb= 172.0)— Quite  recently  (1907),  Urbain  and  Auer  von  Welsbach 
showed  almost  simultaneously  that  the  element  hitherto  regarded  as  ytterbium  is 
a  mixture  of  two  elements,  ytterbium  proper  (neoytterbium),  and  a  new  element 
termed  by  Urbain  lutecium  (At.  wt.  =  174.0). 

Comparison  of  the  Members  of  the  Aluminium 
Group  and  Summary — Corresponding  with  their  position  in 
the  periodic  table,  all  the  members  of  this  group  function  as  tri- 
valent elements.  The  first  member  of  the  boron  sub-group — boron 
itself— is  a  typical  non-metal,  and  the  hydroxide,  B(OH)3,  is  an  acid, 
though  a  comparatively  weak  one.  Aluminium  in  most  of  its  pro- 
perties behaves  as  a  metal,  but,  as  already  indicated,  the  hydroxide, 
A1(OH)3,  is  a  weak  acid  as  well  as  a  weak  base.  Gallium,  indium 
and  thallium,  the  remaining  members  of  the  sub-group,  are  also 
weak  bases  in  their  trivalent  compounds,  as  is  indicated  by  the 
partial  hydrolysis  of  the  chlorides  and  by  other  properties.  It  is 


ELEMENTS   OF   THE   ALUMINIUM   GROUP      491 

only  in  its  trivalent  character,  however,  that  thallium  behaves  like 
other  members  of  the  group.  The  thallous  compounds,  of  the 
type  T1X,  behave  in  some  respects  like  the  alkali  metals,  in  other 
respects  like  silver ;  this  is  shown  by  the  fact  that  thallous  hydroxide 
and  carbonate  are  soluble  in  water,  whereas  the  thallous  halides 
are  comparatively  insoluble.  The  fact  that,  in  contrast  to  thallic 
hydroxide,  thallous  hydroxide  is  a  comparatively  strong  base,  is  of 
considerable  interest  and  importance.  Metallic  thallium  strongly 
resembles  lead. 

The  variation  of  the  properties  of  the  boron  sub-group  with  the 
atomic  weight  is  shown  in  the  accompanying  table: 


B 

Al 

Ga 

In 

Tl 

Atomic  weight    .     . 

II.  0 

27.1 

69.9 

114.8 

204.0 

Density      .... 

2.45 

2.7 

5-9 

74 

11.8 

Melting-point      .     . 
Atomic  volume  .     . 

>2OOO° 
4-5 

657° 

1  0.0 

30.0 
1  1.8 

176° 
15-5 

285° 
17-3 

CHAPTER  XXXI 
ELEMENTS  OF  THE  CARBON  GROUP  (GROUP  IV) 

Sub-group  A  Sub-group  B 

Titanium,  Ti  .        .          48.1  Carbon,  C         .        .         .         12.00 

^Zirconium,  Zr    .         .        .          90.6  Silicon,  Si         .         .        .         28.3 

Cerium,  Ce  .         ,         140.25  Germanium,  Ge        .        .  72.5 

Thorium,  Th      .         .        .        232.4  Tin,  Sn     ....  119.0 

Lead,  Pb  .        .         .  207.1 

THE  sub-group  A  contains  four  rare  elements  ;  and  germanium, 
the  middle  member  of  sub-group  B,  is  also  a  rare  element.  Of 
the  remaining  elements  of  sub-group  B,  carbon  and  silicon  have 
already  been  dealt  with  among  the  non-metals.  Tin  and  lead,  on  the 
other  hand,  are  metals,  but  their  oxides  show  also  weak  acidic  proper- 
ties. In  this,  as  in  certain  other  families  of  elements  already  con- 
sidered, the  electro-positive  character  increases  with  increase  in  atomic 
^veight.  In  accordance  with  their  position  in  the  periodic  table  all  the 
elements  of  this  group  are  quadrivalent,  but  some  of  them,  notably  tin 
and  lead,  show  other  valencies  as  well. 

THE  TITANIUM  SUB-GROUP 

Titanium  (Ti=48.i)  occurs  naturally  as  the  dioxide  in  the  rare  minerals  rutilc, 
anatase  and  brookite,  and  as  ferrous  titanate,  FeTiO3,  in  certain  iron  ores.  The 
element  is  obtained  by  reducing  the  dioxide  with  carbon  in  the  electric  furnace, 
but  does  not  appear  to  have  been  obtained  quite  pure.  It  burns  in  oxygen  to  the 
dioxide,  and  combines  readily  with  nitrogen  at  800°  to  form  the  nitride,  Ti3N4. 
In  its  compounds  it  shows  considerable  analogy  with  silicon  (q.v.  ). 

Titanium  dioxide,  TiO2,  has  both  acidic  and  basic  properties.  When  fused 
with  alkalis  or  alkali  carbonates,  alkali  titanates,  e.g.  K2TiO3,  are  formed.  The 
titanates  are  soluble  in  excess  of  hydrochloric  or  sulphuric  acid,  and  the  solutions 
presumably  contain  titanic  chloride,  TiCl4,  or  sulphate,  Ti(SO4)2. 

When  alkali  hydroxide  or  ammonia  is  added  in  excess  to  either  of  these  solu- 
tions, titanic  acid,  Ti(OH)4,#H2O,  is  obtained  as  a  voluminous  white  precipitate. 
Titanic  chloride,  TiCl4,  is  obtained  as  a  colourless,  fuming  liquid  by  heating  a 
mixture  of  titanium  dioxide  and  carbon,  and  passing  chlorine  over  it.  The 
tetrachloride  is  completely  hydrolyzed  on  addition  of  water  : 


492 


GERMANIUM  493 

Titanium  forms  two  other  series  of  salts,  in  which  it  is  divalent  and  trivalent 
respectively.  The  yellow  colour  obtained  by  the  addition  of  hydrogen  peroxide 
to  a  solution  of  titanium  dioxide  in  concentrated  sulphuric  acid  (p.  143)  is  due  to 
the  formation  of  titanium  peroxide,  TiO3. 

Zirconium  occurs  naturally  chiefly  as  Zircon,  ZrSiO4.  The  metal  is  obtained 
by  reducing  the  dioxide,  ZrO2,  with  carbon  in  the  electric  furnace  under  definite 
conditions.  The  hydroxide,  Zr(OH)4,  is  obtained  as  a  gelatinous  precipitate  by 
adding  ammonia  to  a  solution  of  a  zirconium  salt ;  it  is  only  very  slightly  soluble 
in  aqueous  alkalis,  but  when  fused  with  the  latter  forms  salts  of  the  type  Na2ZrO3, 
which  are  decomposed  by  water.  The  fact  that  zirconium  sulphate,  Zr(SO4)2,  can 
be  recrystallized  from  water  shows  that  zirconium  hydroxide  is  a  base  of  moderate 
strength. 

Cerium  occurs  with  other  elements  (p.  490)  in  the  rare  earths,  more  particularly 
in  cerite,  euxenite  and  monazite.  The  metal,  which  has  been  obtained  by  heat- 
ing the  dioxide  with  magnesium  powder,  has  a  density  of  6.7,  melts  at  623°  and  is 
permanent  in  dry  air.  Cerium  forms  two  well-defined  series  of  salts,  the  cerous 
salts,  e.g.  CeCl3,  in  which  it  is  trivalent,  and  the  eerie  salts,  e.g.  CeCl4,  in  which 
it  is  quadrivalent.  Cerous  oxide,  Ce^O^,  is  obtained  by  heating  the  carbonate  in  a 
current  of  hydrogen.  Ceric  oxide,  CeO2,  is  obtained  as  a  pale  yellow  powder  on 
heating  the  carbonate,  nitrate  or  sulphate  in  air. 

Thorium,  like  cerium,  is  chiefly  obtained  from  monazite  sand  (p.  209),  it  also 
occurs  in  tetragonal  crystals  as  thorite,  ThSiO4,  isomorphous  with  zircon,  rutile 
and  cassiterite,  and  as  thorianite,  a  mineral  obtained  from  Ceylon,  which  is 
chiefly  a  mixture  of  thorium  oxide,  ThO2,  and  uranium  oxide,  UO2.  Thorium 
sulphate,  Th(SO4)2,  and  the  chloride,  ThCl4,  can  be  crystallized  from  aqueous 
solution,  so  that  the  hydroxide,  which  is  insoluble  in  alkalis,  is  a  fairly  strong 
base.  The  dioxide,  ThO2,  is  used  for  the  preparation  of  incandescent  mantles, 
as  already  mentioned.  Thorium  compounds  are  radio-active  (p.  583). 

THE  TIN  SUB-GROUP 

Germaniumi(Ge=72.5),  was  discovered  by  Winkler  (1886)  in  an  argentiferous 
mineral,  argyrodite,  3Ag2S,GeS2  (perhaps  4Ag2S)GeS2),  found  at  Freiberg  in 
Saxony.  It  is  also  found  in  small  proportion  in  some  of  the  rare  earths,  for 
example  in  samarskite  and  in  gadolinite.  The  metal  itself  is  obtained  from  its 
salts  by  heating  in  hydrogen.  It  is  a  brittle,  grayish-white  metal  of  density  5.47 
at  20°.  It  melts  below  900°,  and  is  only  slowly  vaporized  at  1500°.  At  room 
temperature  it  is  not  acted  upon  in  the  air  ;  at  red-heat  it  burns  to  the  dioxide, 
GeO.j.  Germanium  is  insoluble  in  hydrochloric  acid,  but  nitric  acid  converts  it 
into  a  hydrated  dioxide,  GeO2.vH2O. 

Germanium,  like  tin,  forms  two  series  of  salts,  as  it  functions  both  as  a  divalent 
and  tetravalent  element. 

Germaniiim  dichloride,  GeCl2,  is  obtained  by  passing  hydrogen  chloride  over 
the  heated  sulphide,  GeS.  It  has  not  been  much  investigated.  Germanous  sul- 
phide, GeS,  is  obtained  by  heating  a  mixture  of  germanic  sulphide,  GeS2,  and 
germanium  in  a*  current  of  carbon  dioxide.  The  sulphide  occurs  in  grayish-black 
crystals,  and  is  the  best  defined  germanous  compound. 

Germanium  tetrachloride,  GeCl4,  is  obtained  by  heating  germanium  in  a 
current  of  chlorine.  It  is  a  colourless  liquid,  which  boils  at  80°,  and  is  completely 


494    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

decomposed  by  water,  with  formation  of  the  hydroxide,  Ge(OH)4  and  hydro- 
chloric acid. 

Germanium  dioxide,  GeO2,  is  obtained  by  burning  germanium  in  air,  or  by 
heating  the  hydroxide.  It  is  a  white  powder  of  density  4.7  at  78°,  and  is  not 
affected  by  heat.  It  is  soluble  in  water,  and  the  solution  has  an  acid  reaction, 
so  that  the  dioxide  has  acidic  properties,  which  is  confirmed  by  its  ready  solubility 
in  alkalis.  It  is  also  soluble  in  acids,  and  appears  to  possess  slightly  basic 
properties. 

Germanium  hydroxide,  Ge(OH)4,  obtained  as  a  gelatinous  precipitate  by 
hydrolysis  of  the  tetrachloride,  is  a  weak  acid  and  also  a  very  weak  base. 

Germanium  disulphide,  GeS2,  is  obtained  as  a  white  precipitate  by  passing 
hydrogen  sulphide  through  a  solution  of  germanium  dioxide  in  hydrochloric  acid. 
It  is  slightly  soluble  in  water,  readily  soluble  in  solutions  of  alkali  hydroxides  and 
of  alkali  sulphides  (cf,  stannic  sulphide). 

TIN 

Symbol,  Sn.     Atomic  weight,  119.0.     Molecular  weight  uncertain. 

History  —  As  articles  of  bronze  (an  alloy  of  tin  and  copper)  at 
least  4000  years  old  have  been  found  in  Egypt,  it  is  clear  that  the 
metal  must  have  been  known  at  a  very  early  period.  The  ores  used 
by  the  ancient  Egyptians  as  a  source  of  tin  appear  to  have  come  from 
northern  Persia.  In  later  times,  tin  was  obtained,  among  other 
sources,  from  Cornwall  and  Devonshire,  and  it  appears  from  Pliny 
that  the  earliest  name  applied  to  the  British  Islands  was  the  Cassi- 
terides,  from  cassiterite,  the  chief  ore  of  tin.  The  Latin  name  for 
tin  is  Stannum,  hence  the  symbol  Sn. 

Occurrence  —  Tin  does  not  appear  to  occur  naturally  in  the  free 
condition.  It  occurs  almost  exclusively  as  the  dioxide,  SnO2,  cassi- 
terite or  tinstone,  in  tetragonal  crystals,  which  are  usually  coloured 
brown  or  black  by  traces  of  iron.  The  ore  usually  contains  (besides 
iron)  silica,  sulphur,  arsenic,  copper,  lead  and  other  impurities- 
Tinstone  is  found  in  Cornwall,  Saxony,  Bohemia,  Bolivia,  Mexico,  but 
is  now  chiefly  obtained  (generally  of  a  very  high  degree  of  purity) 
from  Banca,  Biliton  and  other  islands  in  the  Straits  Settlements. 

Preparation  of  Metal  —  The  ore  is  first  roasted,  whereby  the 
arsenic  and  sulphur  usually  present  are  removed,  and  is  then  washed 
to  remove  impurities  such  as  copper  sulphate  and  ferric  oxide.  The 
dioxide  is  then  mixed  with  powdered  coal  and  heated  in  a  rever- 
beratory  furnace  : 


The  tin  thus  obtained  is  melted  at  a  low  temperature  and  poured  off 
from  some  of  the  impurities  (the  process  is  termed  liquation),  and  is 


TIN  495 

finally  stirred  with  poles  of  green  wood  (cf.  copper,  p.  425)  in  order  to 
reduce  any  oxide  that  may  have  been  formed. 

Properties  of  Metal  —  Tin  is  a  silvery-white  lustrous  metal  of 
density  7.29  at  15°,  it  melts  at  231.5°,  and  boils  at  2270°.  It  is 
malleable  and  ductile  at  the  ordinary  temperature,  and  is  so  soft  that 
it  can  be  cut  with  a  knife.  When  tin  is  bent  a  peculiar  crackling 
noise,  known  as  the  cry  of  //'«,  is  noticed  ;  it  is  probably  due  to  the 
friction  of  the  crystalline  particles.  The  crystalline  character  of  tin  is 
most  clearly  shown  by  etching  the  surface  with  warm  hydrochloric 
acid  or  aqua  regia.  When  heated  to  200°  tin  becomes  brittle,  and 
can  be  powdered  in  a  mortar. 

Tin  exists  in  different  allotropic  modifications.  At  low  temperatures, 
most  rapidly  at  -  48°,  in  contact  with  an  alcoholic  solution  of  (i  pink 
salt,"  the  ordinary  white  modification  changes  to  a  gray  powder  of 
density  5.8.  E.  Cohen  has  shown  that  the  transition  temperature  for 
the  two  modifications  is  at  18°  (cf.  p.  292)  ;  below  18°  ordinary  tin  in 
contact  with  gray  tin  changes  to  a  powder  of  the  latter  form.  This 
phenomenon  is  called  the  tin  pest.  It  has  sometimes  been  observed 
in  organ-pipes  which  have  been  in  use  for  many  years. 

The  brittle  tin  obtained  by  heating  the  metal  to  200°  is  presumably 
a  third  modification. 

Tin  is  stable  in  the  air  at  room  temperature,  but  on  heating  strongly 
it  burns  to  the  dioxide,  SnO2.  It  is  dissolved  by  hot  hydrochloric 
acid  with  formation  of  stannous  chloride  and  hydrogen  : 

Sn  +  2HCl->SnCI2+H2, 

and  is  also  attacked  by  hot  concentrated  sulphuric  acid,  stannic 
sulphate  and  sulphur  dioxide  being  formed  :  * 


The  action  of  nitric  acid  on  tin  depends  on  the  concentration  of  the 
acid  and  the  temperature.  With  cold  dilute  nitric  acid,  stannous 
nitrate,  Sn(NO3)2,  is  the  chief  product,  but  a  little  stannic  nitrate, 
Sn(NO3)4,  may  also  be  formed.  The  fairly  concentrated  acid  vigor- 
ously attacks  tin,  metastannic  acid  (p.  498)  being  the  chief  product. 
The  strongest  nitric  acid  has  practically  no  action  on  tin. 

Tin   is  dissolved  by  a  boiling  solution   of  sodium   or  potassium 
hydroxide,  with  evolution  of  hydrogen  and  formation  of  alkali  stannate: 


496     A   TEXT-BOOK    OF   INORGANIC    CHEMISTRY 

Uses  of  Tin.  Alloys  —  Tin  is  largely  used  as  a  protective 
coating  for  metals  which  are  acted  on  by  air.  The  ordinary  "tin" 
utensils,  widely  used  for  household  purposes,  are  constructed  of  tin 
plate,  which  is  made  by  dipping  carefully  cleaned  sheets  of  iron  into 
melted  tin  (cj.  p.  435). 

Many  important  alloys  containing  tin  are  in  use.  The  alloys  of 
tin  and  lead  melt  at  a  lower  temperature  than  either  metal  (p.  199). 
Pewter  contains  75  per  cent,  of  tin  and  25  per  cent,  of  lead.  Common 
solder  contains  one  part  of  lead  and  one  part  of  tin  ;  other  solders 
contain  the  metals  in  different  proportions.  Tin  is  an  important 
constituent  of  the  bronzes,  which  have  already  been  referred  to  under 
copper.  Tin  amalgam  is  used  as  a  coating  for  mirrors. 

COMPOUNDS  OF  TIN 

Tin  forms  two  series  of  compounds,  s/annous  compounds,  of  the 
type  SnX2,  in  which  it  is  bivalent,  and  stannic  compounds,  of  the  type 
SnX4,  in  which  the  tin  is  quadrivalent.  The  hydroxides,  Sn(OH)2and 
Sn(OH)4,  are  weak  bases,  and  both  have  acidic  properties. 

STANNOUS  COMPOUNDS 

Stannous  Oxide,  SnO,  and  the  Hydrate,  2SnO,H2O—  Stannous 
oxide  hydrate,  2SnO,H2O,  is  obtained  as  a  white  precipitate  on 
adding  alkali  carbonate  to  a  Stannous  chloride  solution  : 


The  hydrate  is  insoluble  in  ammonium  hydroxide,  but  is  dissolved  by 
sodium  or  potassium  hydroxide,  with  formation  of  an  alkali  stannite. 

2SnO,H2O  +  4NaOH->2Na2SnO2+  3H2O. 

Stannous  oxide  is  obtained  by  heating  the  hydrated  oxide  in  a 
current  of  carbon  dioxide.  It  is  a  black  powder,  which  dissolves  in 
acids  to  form  Stannous  salts.  On  heating  in  air,  it  catches  fire  and 
burns  to  the  dioxide. 

Stannous  Chloride,  SnCl2,  is  obtained  by  dissolving  tin  in 
concentrated  hydrochloric  acid  ;  on  evaporating  the  solution  the 
dihydrate,  SnCl,2H2O,  separates  in  monoclinic  crystals.  The  anhy- 
drous salt  is  obtained  by  heating  pure  tin,  or  Stannous  chloride 
dihydrate,  in  a  current  of  dry  hydrogen  chloride. 

Stannous  chloride  melts  at  250°  and  boils  at  606°.  Vapour  density 
determinations  give  values  for  the  molecular  weight  which  even  at 


TIN  497 

650°  are  not  much  larger  than  those  corresponding  with  the  formula 
SnCl2. 

When  free  from  stannic  salt,  stannous  chloride  dissolves  to  a  clear 
solution  in  a  moderate  amount  of  water,  but  with  excess  of  water, 
a  white  precipitate  of  stannous  oxychloride  is  formed  by  hydrolysis  : 

SnCl2  +  HOH-»Sn(OH)Cl  +  HC1. 

Stannous  chloride  is  a  powerful  reducing  agent,  especially  in  alkaline 
solution,  being  itself  oxidized  to  stannic  hydroxide  or  stannic  chloride. 
It  reduces  mercuric  chloride  to  calomel  and  finally  to  metallic 
mercury  : 


Hg2Cl2  +  SnCl2->2  Hg  +  SnCl4. 

Solutions  of  stannous  chloride  absorb  oxygen  from  the  air,  and 
stannic  hydroxide  is  precipitated  : 

SnCl2  +  O  +  3H2O->Sn(OH)4  +  2HC1. 

In  presence  of  hydrochloric  acid,  stannic  chloride  is  probably  the 
first  product,  but  it  is  ultimately  more  or  less  completely  hydrolyzed 
to  the  hydroxide  : 


+  4HOH-»Sn(OH) 

The  same  equations  express  its  behaviour  with  other  oxidizing  agents. 

Stannous  chloride  is  used  commercially  as  a  mordant  and  as  a 
reducing  agent. 

Stannous  Sulphide,  SnS,  is  described  in  connexion  with 
stannic  sulphide  (see  below). 

STANNIC  SALTS 

Stannic  Oxide,  SnO2,  occurs  naturally  in  crystalline  form  as 
cassiterite.  It  is  prepared  by  burning  tin  in  the  air,  but  is  usually 
obtained  by  strongly  heating  metastannic  acid  (^.-z/.).  As  thus  pre- 
pared, it  is  a  white  amorphous  powder  which  turns  yellow  on  heating, 
but  returns  to  its  original  colour  on  cooling.  It  is  insoluble  in  acids 
or  aqueous  alkalis,  but  when  fused  with  alkalis  or  alkali  carbonates, 
alkali  stannate  is  formed  : 

SnO2  +  2KOH->K2SnO3  +  H2O. 

Hydrates  of  Stannic  Oxide—  At  least  two  hydrates  of  stannic 
oxide,  namely,  stannic  acid,  H2SnO3  (SnO2,H2O),  and  metastannic 

32 


498    A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

acid,  H10Sn6O15  (5[SnO2,H2O]),  appear  to  be  definitely  known.  A 
hydrate  which  corresponds  in  composition  with  orthostannic  acid 
(or  stannic  hydroxide),  Sn(OH)4,  has  also  been  obtained. 

Stannic  Acid,  H2SnO3,  is  obtained  as  a  white  gelatinous  pre- 
cipitate by  the  action  of  ammonium  hydroxide  or  of  calcium  carbonate 
on  stannic  chloride  : 

SnCl4  +  4N  H4OH->SnO2,2H2O  +  4N  H4C1. 

When  dried  in  the  air,  it  has  the  composition  H4SnO4  (  =  SnO2,2H2O)  ; 
when  dried  in  vacuo  or  at  100°,  it  has  the  composition  H2SnO3. 

The  precipitate,  before  drying,  is  soluble  in  sulphuric  and  in 
hydrochloric  acids  ;  the  former  solution  presumably  contains  stannic 
sulphate,  Sn(SO4)2,  which  shows  the  basic  character  of  the  hydroxide. 
The  freshly  precipitated  acid  is  also'  soluble  in  alkalis  with  forma- 
tion of  stannates.  The  compounds  Na2SnO3,3H2O  and  K2SnO3,3H2O 
are  well-defined  salts,  readily  soluble  in  water.  They  can  also  be 
prepared  by  fusing  stannic  oxide  with  the  alkali  hydroxides  or  car- 
bonates. The  sodium  compound  is  used  as  a  mordant  in  dyeing 
under  the  name  of  "  preparing  salt." 

Metastannic  Acid,  H10Sn6O15,  is  obtained  by  the  action  of 
moderately  concentrated  nitric  acid  on  metallic  tin  : 


It  is  probable  that  in  the  first  stage  of  the  reaction  stannic  nitrate 
is  formed  and  is  rapidly  hydrolyzed. 

Metastannic  acid  differs  from  stannic  acid  in  being  insoluble  in  acids. 
With  aqueous  alkalis  it  forms  well-defined  salts,  termed  metastan- 
nates,  e.g.  K2Sn6On,4H2O  (or  K2O55SnO2,4H2O)  and  Na-jSn^On^HgO, 
showing  that  the  acid  is  dibasic.  When  metastannic  acid  is  fused 
with  caustic  alkalis,  alkali  stannates  (not  metastannates)  are  obtained. 
Conversely,  when  stannic  acid  is  heated,  it  partially  changes  to 
metastannic  acid. 

Stannic  Chloride,  SnCl4,  is  prepared  by  passing  chlorine  over 
finely  divided  tin  heated  nearly  to  its  melting-point.  It  is  a  colourless 
liquid  which  boils  at  1  14°  and  fumes  in  the  air.  When  water  is  added 
to  stannic  chloride,  heat  is  given  out  and  a  solid  mass  containing  one 
or  more  hydrates  is  obtained  :  with  more  water  a  clear  solution  is 
formed.  The  solution  is  practically  a  non-conductor  of  electricity 
at  first,  showing  that  Sn<-"  ions  can  only  be  present  in  very  minute 
amount,  but  the  conductivity  gradually  increases,  the  solution  mean- 


TIN 


499 


while  remaining  clear.  It  appears  probable  that  the  chloride  under- 
goes slow  hydrolysis  according  to  the  equation 

SnCl4  +  4HOH->Sn(OH)4  +  4HCl, 

the  stannic  oxide  remaining  as  a  colloidal  hydrosol  (p.  371).  On 
boiling  the  solution,  stannic  acid  is  precipitated. 

Hydrates  of  stannic  chloride  with  3,  4,  5,  8,  and  9  molecules  of 
water  have  been  described.  The  pentahydrate  is  used  as  a  mordant. 

Stannic  chloride  forms  double  salts  with  the  alkali  chlorides,  e.g. 
SnCl4,2H  Cl  and  SnCl4,2N  H4C1.  The  latter,  which  is  known  as  "  pink 
salt,"  is  used  as  a  mordant. 

Tin  Sulphides — Stannous  sulphide,  SnS,  is  made  by  heating 
tin  foil  in  sulphur  vapour  or  by  passing  hydrogen  sulphide  through 
an  acidified  solution  of  stannous  chloride.  It  is  a  dark-brown  powder 
which  is  insoluble  in  solutions  of  the  monosulphides  of  the  alkalis 
(K2S  ;  (NH4)2S),  but  dissolves  in  solutions  of  alkali  polysulphides  (cf. 
p.  519)  with  formation  of  thiostannates : 

SnS  +  (N  H4)2S2-KNH4)2SnS3. 

When  excess  of  hydrochloric  acid  is  added  to  a  thiostannate  solution, 
stannic  sulphide,  SnS2,  is  precipitated : 

(NH4)2SnS3+2HCl->2NH,,Cl  +  SnS2  +  H2S. 

Stannic  Sulphide,  SnS2,  is  obtained  as  above  described  and 
also  by  passing  hydrogen  sulphide  into  a  solution  of  a  stannic  salt. 
Commercially  it  is  prepared  by  heating  in  a  retort  a  mixture  of  tin, 
mercury,  sulphur  and  ammonium  chloride  till  fumes  are  no  longer 
given  off.  The  sulphide  which  remains  in  the  retort  forms  golden- 
yellow  lustrous  scales,  and  is  used  as  a  pigment  under  the  name  of 
mosaic  gold. 

The  thiostannates,  mentioned  above,  bear  the  same  relation  to 
stannic  sulphide  as  the  stannates  do  to  stannic  oxide. 

Tests  for  Tin — All  tin  salts  yield  globules  of  metallic  tin  when 
mixed  with  sodium  carbonate  and  reduced  on  charcoal.  The  re- 
actions of  the  sulphides  of  tin,  as  described  above,  are  characteristic. 
The  mercuric  chloride  test  for  stannous  chloride  is  often  useful. 
Stannic  salts  can  be  converted  to  stannous  salts  by  nascent  hydrogen 
and  the  test  then  applied. 


500     A  TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

LEAD 

Symbol,  Pb.     Atomic  weight,  207.1.     Molecular  weight  unknown. 

History  —  Lead  was  known  to  the  ancient  Egyptians  at  least 
three  thousand  years  ago.  In  ancient  times,  however,  there  was 
some  confusion  between  tin  and  lead  ;  and  Pliny  was  the  first  to 
distinguish  clearly  between  plumbum  nigrum,  lead,  and  plumbum 
album  or  candidun^  tin.  The  metal  was  used  by  the  ancient  Romans 
in  making  water-pipes. 

Occurrence  —  The  chief  ore  of  lead  is  galena,  PbS,  which  is 
very  widely  distributed.  It  also  occurs  in  considerable  quantities 
as  cerussite,  PbCO3  ;  and  in  relatively  small  amount  as  anglesite^ 
PbSO4;  crocoisite,  PbCrO4  ;  ivuljenite,  PbMoO4;  and  pyromorphite, 
PbCl2,3Pb3(P04)2. 

Preparation  of  Metal  —  The  metal  is  obtained  almost  exclu- 
sively from  galena,  which  is  generally  a  very  pure  material.  The  ore 
is  first  roasted  in  a  reverberatory  furnace  (p.  404),  whereby  part  of  the 
sulphide  is  burned  to  oxide,  part  is  oxidized  to  sulphate,  and  a  con- 
siderable proportion  remains  unaltered  : 


2PbS  +  3O2->2PbO  +  2SO2  ; 

The  temperature  is  then  raised,  when  the  sulphate  and  oxide  react 
with  the  sulphide  to  form  the  metal  ; 

PbSO4+  PbS->2Pb  +  2SO2 
2PbO  +  PbS->3Pb  +  SO2. 

In  North  America  and  in  Spain  lead  is  sometimes  obtained  from 
galena  by  heating  with  iron  ;  the  latter  combines  with  the  sulphur  to 
form  ferrous  sulphide,  FeS.  In  practice,  for  economical  reasons,  ores 
which  yield  iron  during  the  process  are  used  instead  of  the  metal  ; 
they  are  mixed  with  galena  and  coke  and  strongly  heated  in  a  blast 
furnace.  The  ferrous  sulphide  impurities  rise  to  the  surface  and  the 
molten  lead  separates  out  below  : 


As  very  pure  lead  is  required  for  many  purposes,  e.g.  for  accumula- 
tors, the  metal  obtained  by  either  of  the  above  processes  must  be 
refined.  Metals  which  are  more  easily  oxidized  than  lead  are  con- 
verted into  oxides  by  heating  the  metal  in  an  oxidizing  atmosphere, 
and  rise  to  the  surface  of  the  melted  lead  as  a  scum.  Other  impuri- 
ties are  removed  by  "poling"  (p.  425). 


LEAD  501 

Properties — Lead  is  a  bluish-white,  very  soft  metal,  very  lustrous 
on  freshly-cut  surfaces  but  rapidly  becoming  dull  owing  to  surface 
oxidation.  It  melts  at  327°  and  boils  at  1525°.  Recent  determina- 
tions of  its  vapour-density  at  1870°  show  that  it  is  monatomic.  Its 
density  at  20°  is  11.34.  Lead  occurs  in  octahedral  crystals  belonging 
to  the  regular  system,  as  can  readily  be  shown  by  melting  it  in  a 
crucible,  allowing  it  partially  to  solidify,  piercing  a  hole  in  the  crust 
and  pouring  off  the  still  liquid  portion. 

From  the  chemical  point  of  view  lead  is  a  fairly  active  metal ;  but 
the  action  of  reagents  on  it  is  in  many  cases  retarded  by  the  formation 
of  protective  coatings  and  by  the  slight  solubility  of  its  compounds. 
The  very  thin  black  film  which  forms  on  the  surface  at  room  temperature 
is  probably  the  suboxide,  Pb2O.  When  lead  is  heated  strongly  in  air 
the  oxide  PbO  is  formed.  The  great  influence  of  the  state  of  division 
on  the  rate  of  reaction  is  well  shown  by  the  fact  that  very  finely 
divided  lead,  obtained  by  reducing  lead  tartrate  in  a  current  of 
hydrogen  at  as  low  a  temperature  as  possible,  catches  fire  in  the  air 
at  room  temperature. 

Lead  is  not  attacked  by  water  free  from  oxygen,  but  water  contain- 
ing oxygen  dissolves  it  with  formation  of  lead  hydroxide  : 

Pb  +  O  +  H2O->Pb(OH)2. 

The  solvent  action  of  natural  waters  on  lead  is  very  important  in  con- 
nexion with  the  extended  use  of  lead  pipes  for  conveying  water. 
Water  saturated  with  air  can  retain  in  solution  more  than  o.  i  mg.  of  lead 
per  litre.  Water  containing  carbonates  or  sulphates,  however,  dissolves 
only  very  minute  quantities  of  lead  owing  to  the  very  slight  solubility 
of  lead  carbonate  and  sulphate  ;  and,  further,  these  salts  form  a  coat- 
ing on  the  metal  which  prevents  further  solvent  action.  Free  carbon 
dioxide  under  certain  circumstances  increases  the  solubility,  and  it  is 
of  advantage  to  add  sufficient  alkali  to  the  water  to  convert  it  to  the 
bicarbonate.  It  follows  from  the  above  that  lead  pipes  can  be  used 
safely  for  most  natural  waters ;  but  very  pure  waters,  e.g.  rain  water, 
may  take  up  dangerous  amounts  of  lead.  The  water  of  Loch  Katrine, 
used  by  the  city  of  Glasgow,  is  filtered  through  beds  of  chalk,  and 
the  carbonate  thus  taken  up  greatly  lessens  the  solvent  action  of 
the  water  on  the  lead  pipes.1  Water  containing  ammonium  salts,  and 

1  The  main  object  in  filtering  the  water  through  chalk  was  to  supply  the  lime 
salts  essential  for  the  proper  nourishment  of  the  human  organism.  It  is  now 
held,  however,  that  sufficient  lime  for  this  purpose  is  present  in  ordinary  foods 
(especially  farinaceous  foods), 


502     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

also  water  containing  weak  organic  acids,  e.g.  acetic  acid,  have  con- 
siderable solvent  action  on  lead. 

Lead  is  rapidly  dissolved  by  nitric  acid.  Hydrochloric  acid  is 
practically  without  action  at  room  temperature  ;  but  the  hot  concen- 
trated acid  acts  slowly  on  it,  with  formation  of  lead  chloride  and 
hydrogen.  Hot  concentrated  sulphuric  acid  slowly  dissolves  lead  as 
the  sulphate,  PbSO4  ;  but  a  less  concentrated  acid  is  practically  with- 
out action,  owing  to  the  insolubility  of  lead  sulphate. 

Lead,  being  a  weakly  electro-positive  metal,  is  displaced  from  its 
salts  in  solution  by  zinc,  magnesium,  and  other  metals  (p.  434).  When 
a  strip  of  zinc  is  suspended  in  a  dilute  solution  of  lead  acetate,  the 
lead  is  deposited  as  a  branching  crystalline  structure  known  as  the 
lead  tree  : 

Pb(C2H302)2  +  Zn->Zn(C2H302)2  +  Pb. 

Uses  of  Lead.  Alloys — On  account  of  the  readiness  with 
which  it  can  be  worked  and  its  resistance  to  many  reagents  lead  is 
largely  used  for  commercial  purposes.  Lead  pipes  are  made  by  heat- 
ing the  metal  till  soft  and  squeezing  it  into  shape  by  hydraulic  pressure. 
Lead  bullets,  which  contain  a  small  proportion  of  arsenic,  are  made 
by  forcing  the  metal  into  moulds.  Alloys  of  lead  with  tin  have 
already  been  referred  to  (p.  496) ;  lead-antimony  and  lead-bismuth 
alloys  will  be  mentioned  at  a  later  stage. 

COMPOUNDS  OF  LEAD 

Lead,  like  tin,  forms  two  series  of  compounds,  in  which  it  functions 
as  a  divalent  and  as  a  quadrivalent  element.  In  its  divalent  com- 
pounds, which  are  by  far  the  more  important,  it  is  a  base  of  moderate 
strength,  and  has  very  weak  acidic  properties  ;  in  its  tetravalent 
compounds  it  has  scarcely  any  basic  properties,  the  dioxide,  PbO2, 
being  distinctly  acidic.  The  compounds  containing  divalent  lead 
are  usually  called  plumbic  compounds;  the  compounds  containing 
quadrivalent  lead  have  no  general  designation. 

Oxides  of  Lead— Five  oxides  of  lead  are  known  :  Pb2O,  PbO, 
Pb2O3,  Pb3O4,  and  PbO2. 

Lead  Suboxide,  Pb2O,  is  obtained  by  heating  lead  oxalate  in  a 
current  of  carbon  dioxide  or  nitrogen  at  the  lowest  temperature  suffi- 
cient to  effect  decomposition  (about  300°)  : 

2PbC2O4-»Pb2O  +  3CO2  +  CO. 

It  is  also  formed  by  the  action  of  air  on  lead  at  temperatures  below 
its  melting-point. 


LEAD  503 

Lead  suboxide  is  a  grayish-black  powder,  which  is  decomposed  on 
heating  into  plumbic  oxide,  PbO,  and  lead.  Acids  act  on  it  in  an 
analogous  way  ;  a  plumbic  salt  goes  into  solution  and  lead  remains  : 


Plumbic  Oxide  (Litharge},  PbO,  is  obtained  as  a  yellowish-red 
amorphous  powder  by  strongly  heating  lead  in  air,  and  is  a  by-pro- 
duct in  the  separation  of  lead  from  silver  by  cupellation  (p.  436).  It  is 
also  obtained  by  heating  the  nitrate  or  carbonate,  and  also  by  heating 
any  of  the  other  oxides  at  a  sufficiently  high  temperature  in  air. 
When  lead  oxide  is  heated  above  its  melting-point,  835°,  and  allowed 
to  cool  it  solidifies  to  a  crystalline  mass  known  as  litharge.  There 
are  at  least  two  modifications  of  lead  oxide,  a  yellow  and  a  red,  the 
latter  being  the  stable  form  at  room  temperature. 

Lead  oxide  is  slightly  soluble  in  water,  forming  an  alkaline  solution 
which  presumably  contains  lead  hydroxide,  Pb(OH)2.  It  dissolves 
in  acids  to  form  plumbic  salts,  and  also  dissolves  on  boiling  with 
alkali  hydroxides,  owing  to  the  formation  of  soluble  plumbites,  e.g. 
Na2PbO2  (see  below). 

Lead  Hydroxide,  Pb(OH)2,  is  formed  by  the  slow  oxidation 
of  lead  in  moist  air,  and  also  as  a  white  precipitate,  by  adding  an  alkali 
hydroxide  or  ammonium  hydroxide  to  the  aqueous  solution  of  a  lead 
salt.  The  nature  of  the  precipitate  appears  to  depend  on  the  condi- 
tions ;  the  hydrates  2PbO,H2O  and  3PbO,H2O  have  been  described. 

The  basic  character  of  lead  hydroxide  is  shown  by  the  fact  that  it 
dissolves  to  a  small  extent  in  water,  forming  an  alkaline  solution,  and 
reacts  with  acids  to  form  corresponding  salts.  Its  acidic  character  is 
shown  by  its  solubility  in  solutions  of  alkali  hydroxides;  the  solutions 
contain  plumbites,  e.g.  Na2PbO2  : 

Pb(OH)2  +  2NaOH->Pb(ONa)2  +  2H2O. 

Lead  Sesquioxide,  Pb2O3,  is  obtained  as  an  orange-yellow 
powder  by  adding  sodium  hypochlorite  to  a  solution  of  plumbic  oxide 
in  alkali.  Acids  decompose  it  into  the  monoxide  and  dioxide,  the 
former  of  which  dissolves  to  form  plumbic  salts  ;  and  it  may  there- 
fore be  regarded  as  a  loose  compound  of  the  two  oxides,  PbO  and 
PbO2  (cf.  red  lead). 

Red  Lead  or  Minium,  Pb3O4,  is  obtained  by  prolonged  heat- 
ing of  lead  monoxide  in  the  air  at  500°.  It  is  a  bright  scarlet  powder, 
and  is  used  as  a  pigment.  When  heated  to  550°  the  dissocia- 
tion pressure  of  the  oxygen  is  equal  to  the  oxygen  pressure  in  the 


504     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

atmosphere,  and  the  compound  therefore  changes  to  the  monoxide. 
With  acids  the  monoxide  is  dissolved  out  with  formation  of  plumbic 
salts,  and  lead  dioxide  remains  : 


so  that  the  salt  behaves  as  a  mixture  of  the  two  oxides.  It  has, 
however,  been  shown  that  a  definite  compound  of  the  formula 
Pb3O4(PbO2,2PbO)  exists,  since  the  dissociation  pressure  of  the  oxygen 
in  red  lead  is  less  than  that  of  lead  peroxide  at  the  same  temperature. 
The  composition  of  the  commercial  article  is  not  quite  constant  ; 
besides  the  compound,  Pb3O4,  it  may  contain  excess  of  PbO  or 
PbO2. 

Lead  Dioxide,  PbO2,  is  obtained,  as  just  described,  by  the 
action  of  dilute  nitric  acid  on  red  lead  ;  it  is  also  obtained  by  the 
oxidation  of  lead  monoxide,  dissolved  in  alkali,  with  hypochlorite,  and 
is  deposited  on  the  anode  during  the  electrolysis  of  lead  salts  (cf. 
p.  507).  It  is  an  amorphous,  dark-brown  powder,  which  on  heating 
to  350°  gives  up  oxygen  with  formation  of  the  lower  oxides  mentioned 
above  ;  when  the  temperature  is  sufficiently  high  the  monoxide 
results.  When  heated  with  concentrated  hydrochloric  acid  chlorine 
is  given  off: 


With  concentrated  sulphuric  acid  oxygen  is  given  off  : 


In  both  cases  plumbic  salts  are  formed. 

Lead  dioxide  shows  both  acidic  and  basic  characters.  When 
boiled  with  a  concentrated  solution  of  potassium  hydroxide  it  dis- 
solves, and  from  the  solution,  on  cooling,  potassium  plumbate, 
K2PbO3,3H2O,  separates  in  colourless  rhombohedral  crystals,  iso- 
morphous  with  potassium  stannate  trihydrate  (p.  498.).  Salts  corre- 
sponding with  "  orthoplumbic  acid,"  Pb(OH)4,  are  also  known,  e.g. 
Ca2PbO4.  This  salt  is  obtained  by  heating  in  the  air  a  mixture  of 
calcium  carbonate  and  lead  oxide,  oxygen  being  absorbed  : 


The  action  is  reversible,  and  upon  this  is  based  Kassner's  method  (no 
longer  used)  of  obtaining  oxygen  from  the  air.  The  compound 
Pb3O4  may  be  regarded  as  plumbic  plumbate,  Pb2PbO4. 

Lead  Chloride,  PbCl2,  being  only  slightly   soluble   in   water, 
separates  when  solutions  containing  Pb"  and  Cl'  ions  are  brought 


LEAD  5°5 

together.  It  forms  colourless,  lustrous  rhombic  crystals  ;  at  20°  100 
grams  of  water  dissolve  about  i.o  gram,  at  100°  about  4  grams  of  the 
salt.  The  solubility  is  diminished  by  the  addition  of  dilute  hydro- 
chloric acid  (and  soluble  chlorides),  but  is  greater  in  the  presence  of 
concentrated  hydrochloric  acid.  The  lowering  of  solubility  is  due  to 
the  addition  of  a  compound  with  a  common  ion  (p.  440)  ;  the  subse- 
quent increase  to  the  formation  of  complex  ions  (p.  440).  In  the 
latter  case,  the  solutions  probably  contain  compounds  of  the  type, 
HPbCl3,  which  ionise  as  follows: 


Lead  Bromide,  PbBr2,  and  Lead  Iodide,  PbI2,  are  prepared 
in  a  similar  manner  to  the  chloride  ;  the  former  is  colourless,  the 
latter  separates  from  aqueous  solution  in  lustrous  yellow  crystals.  At 
o°  100  grams  of  water  dissolves  0.044  grams,  at  25°  0.076  grams,  and 
at  100°  0.436  grams  of  the  iodide.  All  the  lead  halides  form  double  salts 
with  the  alkali  halides,^.  2PbCl2-KCl  ;  PbCl2'2KCl;  PbI2'KI'2H2O. 
Solutions  of  these  double  salts  contain  the  lead,  in  part  at  least,  as  a 
complex  anion,  e.g.  PbCl3'  (see  above). 

Lead  Nitrate,  Pb(NO3)2,  is  obtained  by  dissolving  lead,  the 
oxide  or  carbonate  in  nitric  acid  ;  on  cooling  it  separates  in  anhydrous 
octahedral  crystals.  At  o°  100  grams  of  water  dissolve  39  grams,  at 
20°  56  grams  of  the  salt.  On  heating  it  is  decomposed  into  lead  oxide, 
nitrogen  peroxide  and  oxygen  (p.  228). 

A  number  of  basic  nitrates  of  lead,  e.g.  Pb(NO3)2'2PbO  and 
Pb(NO3)2'PbO,  have  been  described. 

Lead  Sulphate,  PbSO4,  occurs  naturally  in  rhombic  crystals  as 
anglesite^  and  is  formed  when  solutions  containing  Pb"  and  SO4" 
ions  are  mixed.  It  is  a  heavy  white  powder,  practically  insoluble  in 
water,  still  less  soluble  in  dilute  sulphuric  acid,  but  readily  dissolves 
in  concentrated  sulphuric  acid.  The  latter  phenomenon  is  due  to  the 
formation  of  an  acid  sulphate,  Pb(HSO4)2,H2O,  which  separates  when 
the  solution  in  question  is  cautiously  diluted.  Several  basic  sulphates 
of  lead  have  been  described. 

Lead  sulphate  dissolves  fairly  readily  in  solutions  of  ammonium 
acetate  and  of  the  alkali  acetates.  In  dilute  solution  this  effect  is  due 
to  the  formation  of  slightly  ionised  lead  acetate  by  double  decomposi- 
tion and  a  consequent  diminution  in  the  Pb"  ion  concentration.  In 
concentrated  solution  the  formation  of  complex  salts  appears  to  come 
into  account. 

Lead  Carbonate,  PbCO3,  occurs  naturally  as  cerussite.     It  is 


5o6     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

obtained  as  a  heavy  white  precipitate  by  adding  a  solution  of  lead 
nitrate  or  acetate  to  excess  of  a  solution  of  commercial  ammonium 
carbonate.  When  sodium  or  potassium  carbonate  is  used,  basic  car- 
bonates of  lead,  ^PbCO3,jPb(OH)2,  are  obtained,  the  composition  of 
which  depends  upon  the  conditions  of  precipitation. 

The  best  known  basic  carbonate  of  lead,  2PbCO3,Pb(OH)2,  known 
as  white  lead,  is  largely  used  as  a  pigment.  A  number  of  methods  for 
preparing  it  are  in  use,  but  the  product  of  highest  covering  power  is 
obtained  by  the  so-called  Dutch  process,  which  has  been  used  for 
centuries.  The  method  depends  upon  the  successive  action  of  acetic 
acid  and  of  carbon  dioxide  on  the  metal.  The  lead  is  cast  in  gratings, 
or  sheet  lead  is  twisted  in  spirals  so  as  to  expose  a  large  surface,  and 
the  gratings  or  spirals  are  placed  in  earthware  pots  in  such  a  way  that 
they  do  not  come  into  contact  with  the  small  quantity  of  vinegar 
(dilute  acetic  acid)  which  each  pot  contains.  A  row  of  pots  is  placed 
on  dung  or  spent  tan  bark  in  a  shed,  and  alternate  rows  of  pots  and 
layers  of  tan-bark  or  manure  are  piled  up  till  the  shed  is  nearly  full. 
The  whole  arrangement  is  left  for  some  weeks.  The  heat  given  out 
during  the  fermentation  of  the  tan-bark  vaporizes  the  acetic  acid, 
which,  along  with  oxygen,  converts  the  lead  to  a  basic  acetate, 
Pb(C2H3O2)2,Pb(OH)2.  The  latter  is  then  in  turn  acted  upon  by  the 
carbon  dioxide  produced  during  the  fermentation,  the  basic  carbonate 
being  formed.  When  the  gratings  are  almost  entirely  converted  to 
white  lead,  the  latter  is  scraped  off,  ground  up  while  wet,  washed  to 
remove  lead  acetate,  and  dried. 

White  lead  is  also  obtained  by  passing  carbon  dioxide  into  a  solu- 
tion of  basic  lead  acetate,  or  by  rubbing  lead  oxide  into  a  paste  with 
lead  acetate  solution  and  then  passing  in  carbon  dioxide.  Electrolytic 
methods  have  also  been  proposed.  None  of  these  methods  gives  a 
product  equal  to  that  obtained  by  the  Dutch  process. 

White  lead  has  the  disadvantage  that  it  is  poisonous,  and  it  blackens 
when  exposed  to  hydrogen  sulphide,  which  is  usually  present  in  the 
atmosphere  of  towns,  but  these  drawbacks  are  more  than  counter- 
balanced by  its  great  covering  power.  In  order  to  give  good  results 
it  must  be  amorphous. 

Lead  Sulphide,  PbS,  occurs  naturally  as  galena,  and  is  obtained 
as  a  black  amorphous  precipitate  by  passing  hydrogen  sulphide  into  a 
solution  of  a  lead  salt.  As  the  sulphide  is  practically  insoluble  in 
water,  hydrogen  sulphide  is  used  to  detect  minute  amounts  of  lead  in 
water.  Lead  sulphide  is  insoluble  in  dilute  hydrochloric  acid,  but  the 
concentrated  acid  converts  it  to  lead  chloride  with  evolution  of 


LEAD  507 

hydrogen  sulphide.  When  lead  sulphide  is  boiled  with  dilute  nitric 
acid,  lead  nitrate  is  formed  ;  the  concentrated  acid  converts  it  chiefly 
into  the  sulphate. 

When  hydrogen  sulphide  is  passed  into  a  solution  of  lead  chloride, 
a  reddish  precipitate  is  obtained  which  becomes  black  when  excess  of 
hydrogen  sulphide  is  used.  This  phenomenon  is  due  to  the  inter- 
mediate formation  of  one  or  more  double  compounds  of  lead  sulphide 
and  chloride.  The  only  one  which  has  been  definitely  isolated  has 
the  formula  PbSfPbCl2  ;  it  is  a  red  powder. 

Lead  Acetate,  Pb(C2H3O2)2,3H2O,  known  as  "  sugar  of  lead  " 
on  account  of  its  sweetish  taste,  is  obtained  in  monoclinic  crystals 
by  dissolving  lead  oxide  in  acetic  acid  and  concentrating  the 
solution.  It  is  very  soluble  in  water.  With  lead  oxide  it  forms 
soluble  basic  salts  ;  two  such  compounds,  Pb(C2H3O2)2,Pb(OH)2  and 
Pb(C2H3O2)2,2Pb(OH)2,  have  been  definitely  isolated. 

Compounds  of  Quadrivalent  Lead— Mention  has  already  been  made  of  lead 
dioxide,  PbO2,  and  of  salts  derived  from  the  corresponding  hydroxide,  Pb(OH)4, 
functioning  as  orthoplumbic  acid,  and  from  its  first  anhydride,  H2PbO3,  meta- 
plumbic  acid.  The  compounds  derived  from  the  hydroxide,  Pb(OH)4,  acting  as 
a  base  will  now  be  dealt  with. 

Lead  Tetrachloride,  PbCl4,  is  formed  when  lead  dioxide  is  dissolved  in  cold 
concentrated  h/drochloric  acid.  It  is  most  readily  obtained  by  passing  chlorine 
into  water  in  which  lead  dichloride  is  suspended,  and  then  adding  ammonium 
chloride  to  the  solution,  when  the  compound,  PbCl4,2NH4Cl,  separates  in 
crystalline  form.  The  latter  compound  is  added  to  cooled  concentrated  sulphuric 
acid,  when  the  tetrachloride  separates  as  a  yellow,  highly  refractive  liquid. 

Lead  tetrachloride  is  a  heavy  yellow  liquid  of  density  3. 18  at  o°.  With  a  small 
amount  of  water  it  forms  a  hydrate  ;  with  excess  of  water  it  is  completely  hydro- 
lyzed  to  the  dioxide  and  hydrochloric  acid : 

PbCl4+4HOH->PbO2+4HCl  +  2H2O. 

Lead  Disulphate,  Pb(SO4)2,  is  obtained  by  the  electrolysis  of  sulphuric  acid  of 
density  1.7  to  1.8  between  electrodes  of  lead,  the  temperature  not  being  allowed 
to  rise  above  30°.  Towards  the  end  of  the  electrolysis  the  temperature  is  raised  to 
40  to  50°,  and  on  cooling  the  disulphate  separates  from  the  solution  in  the  anode 
compartment  in  yellowish  crystals.  When  pure  it  is  colourless. 

Lead  disulphate  is  readily  decomposed  by  water  into  the  dioxide  and  sulphuric 
acid,  a  basic  salt,  PbOSO4,H2O,  being  formed  as  intermediate  product.  It  is  a 
powerful  oxidizing  agent.  Lead  tetracetate,  Pb(C2H3O2)4,  has  also  been 
obtained  ;  it  occurs  in  colourless  needles. 

The  Lead  Accumulator 1 — The  lead  accumulator  is  a  cell  in 
which  electrical  energy  is  stored.     It  consists  in  its  simplest  form  of 
two  plates,  one  of  which  when  charged  is  coated  with  lead  peroxide,  the 
1  Cf.  Physical  Chemistry,  p.  400. 


508    A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 


other  with  finely-divided  metallic  lead,  and  the  plates  dip  into  dilute 
sulphuric  acid.  When  the  two  poles  are  connected  by  a  wire,  dis- 
charge takes  place  at  a  potential  of  two  velts,  and  the  potential 
remains  practically  constant  for  a  long  time.  During  the  discharge, 
the  lead  dioxide  becomes  reduced  to  the  oxide,  which  forms  lead 
sulphate  with  the  sulphuric  acid  ;  at  the  other  pole  the  lead  is 
oxidized  to  lead  sulphate.  The  chemical  changes  taking  place  when 
the  cell  is  discharging  are  therefore  represented  by  the  equation  : 


The  essential  feature  of  the  accumulator  is  that  it  is  readily  rever- 
sible. When,  after  it  is  discharged,  a  current  is  sent  through  it  in 
the  opposite  direction  to  the  discharge  current,  the  sulphate  at  one 
pole  is  oxidized  to  the  dioxide,  at  the  other  pole  it  is  reduced  to 
metallic  lead,  the  cell  being  thus  restored  to  the  same  condition  as 
before  discharge. 

Tests  for  Lead  —  Lead  compounds  are  readily  reduced  on 
charcoal  to  the  metal,  which  can  be  recognized  by  its  softness  and 
by  its  property  of  marking  paper.  The  formation  and  behaviour 
of  the  sulphide,  sulphate  and  iodide,  already  fully  described,  are 
characteristic. 

General  Properties  of  the  Carbon  Sub-group  and 
Summary  —  Carbon,  silicon,  germanium,  tin  and  lead  constitute 
a  natural  family  of  elements  and  show  the  usual  gradual  change  of 
physical  and  chemical  properties  with  increase  of  atomic  weight. 
The  more  important  physical  properties  are  summarized  in  the 
accompanying  table  : 


Carbon. 

Silicon. 

Germanium. 

Tin. 

Lead. 

Atomic  weight 

12.00 

28.3 

72-5 

119.0 

207.1 

Density 
Melting-point 

2.25  to  3.6 
very  high 

2.35  to  2.49 
1600° 

5-47 

<9oo° 

7.29 
231-5 

ii-34 

334° 

Boiling-point 

3000°? 

>i5oo° 

2270° 

1525° 

Atomic  volumes 

3-4 

II.O 

13.2 

16.6 

'18.3 

As  regards  the  chemical  properties,  all  the  elements  are  tetravalent, 
corresponding  with  their  position  in  the  periodic  table,  and  the  last 
three  form  well-defined  divalent  compounds.  Carbon  is  also  divalent 
in  carbon  monoxide.  Further,  the  acidic  character  diminishes  with 
increase  of  atomic  weight.  Carbon  and  silicon  are  typical  non- 


. 
ELEMENTS   OF   THE   CARBON   GROUP        509 

metals ;  no  compounds  are  known  containing  carbon  or  silicon 
cations.  The  oxides  of  germanium,  tin  and  lead  are  both  basic  and 
acidic,  and  this  applies  to  the  oxides  in  which  the  metals  are  divalent 
as  well  as  to  those  in  which  they  are  quadrivalent.  All  the  hydroxides 
of  the  type  X(OH)4  are,  however,  very  weak  bases  (including  H4PbO4), 
and  the  halogen  compounds  of  the  type  XCl4are  immediately  decom- 
posed by  water.1  In  the  case  of  the  divalent  compounds  there  is  a 
definite  increase  of  electro-positive  character  with  increase  of  atomic 
weight.  Stannous  hydroxide,  Sn(OH)2,  is  a  fairly  strong  base,  and 
stannous  salts  with  strong  acids  are  not  greatly  hydrolyzed  in  solution. 
Lead  hydroxide,  Pb(OH)2,  is  a  stronger  base  than  stannous  hydroxide, 
and  plumbic  salts  with  strong  acids,  e.g.  lead  nitrate,  are  scarcely 
hydrolyzed  in  aqueous  solution  at  moderate  dilution. 

In  accordance  with  the  gradation  of  properties  in  the  group,  only 
the  first  two  elements  form  compounds  with  hydrogen. 

1  Carbon  tetrachloride,  CC14,  is  an  apparent  exception  to  this  rule,  as  it  is  not 
affected  by  water  in  the  cold.  It  is,  however,  completely  decomposed  by  heating 
with  water  in  a  sealed  tube,  and  the  action  is  not  reversible : 


s 


CHAPTER  XXXII 
ELEMENTS  OF  THE  NITROGEN  GROUP  (GROUP  V) 

Sub-group  A  Sub-group  B 

Vanadium,  V.         .         .         .       51.06       Nitrogen,  N    ....  14.01 

Columbium,  Cb  (Niobium)     .       93.5          Phosphorus,  P         .         .         .  31.0 

Tantalum,  Ta                                181.0         Arsenic,  As      ....  74.96 

Antimony,  Sb.        .         .         .  120.2 

Bismuth,  Bi    .         .         .         .  208.0 

THE  three  members  of  sub-group  A  are  rare  elements.  Of  the 
members  of  sub-group  B,  nitrogen  and  phosphorus  have  already 
been  considered  among  the  non-metals.  This  family  presents  a  very 
striking  illustration  of  the  increase  in  electro-positive  character  with 
increase  of  atomic  weight.  Nitrogen  and  phosphorus  are  typical 
non-metals,  arsenic  behaves  in  practically  all  respects  like  a  non- 
metal,  antimony  is  intermediate  in  character,  bismuth  behaves  in 
most  respects  like  a  metal.  Throughout  the  group  the  principal 
valencies  are  three  and  five. 

THE  VANADIUM  SUB-GROUP 

Vanadium  was  discovered  by  del  Rio  in  1801  in  vanadinite,  3Pb3(VO4)2'PbCl2, 
isomorphous  with  pyromorphite  (p.  500),  which  is  found  chiefly  in  Spain,  Chili, 
and  the  Argentine.  The  metal  was  first  obtained  by  Roscoe  (1867)  by  heating 
the  dichloride,  VC12,  in  a  current  of  hydrogen,  but  was  not  pure.  The  pure 
metal  was  obtained  for  the  first  time  (Weiss  and  Aichel,  1904)  by  the  reduction 
'of  vanadium  pentoxide  by  the  thermite  process  (p.  482),  a  mixture  of  rare  earth 
metals  being  used  in  place  of  aluminium.  Methods  for  preparing  vanadium  by 
electrolysis  have  also  been  proposed. 

Vanadium  is  a  white,  lustrous,  crystalline  metal  of  density  about  5.8  ;  it  is  the 
hardest  metal  known,  is  fairly  brittle,  and  melts  about  1680°.  It  is  stable  in  the 
air  at  room  temperature,  but  on  heating  burns  to  V2O5. 

Vanadium  is  remarkable  for  the  numerous  series  of  compounds  derived  from 
it.  There  are  four  well-defined  oxides,  VO  (or  V2O2),  V2O3,  V2O4,  and  V2O5. 
These  compounds  are  of  analogous  type  to  the  oxides  of  nitrogen,  but  as  regards 
the  lower  compounds  there  is  very  little  resemblance  in  their  chemical  properties. 
VO  is  a  strongly  basic  oxide,  and  well-defined  salts  corresponding  with  it  are 
known.  The  sulphide,  VSO4,7H2O,  forms  reddish-violet  crystals,  isomorphous 
with  ferrous  sulphate  heptahydrate,  FeSQ^HoO.  Vanadous  chloride,  VC12, 
occurs  in  greenish  hexagonal  plates.  The  oxide  V2O3  is  a  base  of  moderate 

510 


ARSENIC  511 

strength  ;  VClj,6H2O  occurs  in  green  crystals,  V2(SO4)3  is  a  yellow  powder 
(cf.  ccm pounds  of  trivalent  iron).  The  oxide  VO2  is  weakly  basic  and  also 
weakly  acidic.  The  tetrachloritie,  VC14,  is  known,  but  most  of  the  compounds 
containing  quadrivalent  vanadium  are  of  the  type  VOX2  ;  that  is,  in  solution 
divalent  VO"  ions  are  present.  The  pentoxide,  V2O5,  is  acidic,  and  behaves  in 
many  respects  like  phosphorus  pentoxide.  The  salts  are  derived  from  ortho- 
vanadic  acid,  H3VO4,  and  metavanadic  acid,  HVO3 ;  the  latter  are  the  more 
stable.  Ammonium  metavanadate,  NH4VO8,  is  insoluble  in  ammonium  chloride 
solution,  and  advantage  is  taken  of  this  in  separating  vanadium  from  its  ore. 

Salts  containing  di-  and  trivalent  vanadium  are  powerful  reducing  agents. 

Vanadium  is  now  finding  considerable  application  as  a  constituent  of  certain 
alloys. 

ColumM'um  and  Tantalum  occur  together  in  the  rare  minerals  columbite 
and  tantalite,  Fe[Cb(Ta)O3]2.  Both  elements  are  always  present;  when 
columbium  is  in  excess  it  is  termed  columbite,  with  excess  of  tantalum 
tantalite.  Niobium  is  a  lustrous,  fairly  ductile  metal  which,  when  pure,  is 
scarcely  as  hard  as  soft  steel,  but  when  traces  of  impurities  are  present,  is  nearly 
as  hard  as  vanadium  ;  it  melts  about  1950°.  Metallic  tantalum  can  be  obtained 
by  reduction  of  the  dioxide,  TaO2,  with  carbon  in  the  electric  furnace ;  the  powder 
thus  obtained  is  purified  and  obtained  in  coherent  form  by  heating  in  the  electric 
arc  in  a  vacuum,  when  the  oxide  distils  off,  leaving  the  fused  metal,  which  solidifies 
on  cooling  to  a  lustrous  regulus. 

Pure  tantalum  is  a  lustrous  metal  of  grayish  colour,  and  is  so  ductile  that  it 
can  be  beaten  into  foil  and  drawn  out  into  wire  ;  it  melts  about  2275°.  At  a 
white  heat,  it  burns  slowly  in  air  to  the  pentoxide.  As  is  well  known,  tantalum 
is  now  largely  used  as  a  filament  for  electric  lamps. 

The  pentoxides  Cb2O5  and  Ta^Og  are  both  basic  and  acidic.  The  salts  derived 
from  niobic  and  tantalic  acids  are  of  several  types  ;  their  behaviour  shows  that 
the  acids  are  weaker  than  vanadic  acid.  The  remarkably  close  resemblance  be- 
tween niobium  and  tantalum  compounds  is  further  shown  by  the  fact  that  there 
is  no  definite  difference  in  strength  between  niobic  and  tantalic  acids. 

Niobium  and  tantalum  differ  markedly  from  vanadium  in  showing  very  little 
tendency  to  form  compounds  of  a  low  stage  of  oxidation.  Only  one  compound 
of  the  type  CbX3  is  known  with  certainty,  and  tantalum  is  exclusively  quadrivalent 
and  quinquevalent  in  its  compounds. 


THE  NITROGEN  SUB-GROUP 
N  =  14.01.     P  =  3i'.o.     As  =  74.96.     Sb=  120.2.     Bi  =  208.0. 

Of  the  members  of  this  family  only  arsenic,  antimony,  and  bismuth 
remain  to  be  dealt  with. 

ARSENIC 

Symbol,  As.     Atomic  weight=75.     Molecular  weight  =  300. 

Occurrence — Arsenic  occurs  free  in  nature,  but  is  usually  met 
with  in  combination.  In  association  with  oxygen  it  occurs  as  arseno- 
life,  As2O3,  and  with  sulphur  as  realgar,  As2S2,  and  orpiment,  As2S3. 


512     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

In  combination  with  metals  it  occurs  as  arsenical  iron,  FeAs2, 
arsenical  nickel  or  kupfernickel,  NiAs,  and  cobalt  speiss,  CoAs2. 
In  association  with  metals  and  sulphur  it  occurs  in  arsenical  pyrites, 
FeSAs  or  FeS2,FeAs2,  in  cobalt  glance,  CoSAs,  and  in  other  com- 
pounds. It  occurs  in  most  specimens  of  iron  pyrites,  and  hence  finds 
its  way  into  commercial  sulphuric  acid  and  into  products,  such  as 
phosphorus,  into  the  preparation  of  which  sulphuric  acid  enters. 

Preparation  of  Element — (i)  On  the  commercial  scale  it  is 
obtained  by  sublimation  of  native  arsenic  or  by  heating  arsenical 
pyrites  in  earthenware  vessels  in  absence  of  air.  The  pyrites  are 
decomposed  into  ferrous  sulphide  and  arsenic,  and  the  vapour  of  the 
latter  is  condensed  in  earthenware  receivers : 

FeS2,FeAs2-»2FeS  +  2As. 

It  is  purified  by  sublimation. 
(2)  It  is  easily  obtained  by  heating  arsenious  oxide  with  charcoal : 

As2O3+3C->2As-t-3CO. 

Properties — Arsenic  exists  in  at  least  three  allotropic  modifica- 
tions. The  best  known  form  (metallic  arsenic)  occurs  in  rhombo- 
hedral  crystals,  which  are  steel-gray  in  colour,  and  have  metallic 
lustre  ;  the  density  is  5.73.  Metallic  arsenic  is  brittle  ;  it  is  a  good 
conductor  of  heat  and  electricity.  It  sublimes  at  450°,  before  the 
melting-point  is  reached,  but  can  be  melted  under  pressure.  A 
second  form  is  obtained  as  a  black,  apparently  amorphous,  but  really 
minutely  crystalline,  powder  by  subliming  ordinary  arsenic  ;  its  density 
is  4.72.  At  300°  it  changes  with  evolution  of  heat  into  the  stable 
metallic  form.  A  third  modification  is  obtained  in  light  yellow 
crystals  (density  2.026)  by  rapidly  cooling  arsenic  vapour,  e.g.  by 
means  of  liquid  air.  This  modification,  like  yellow  phosphorus,  is 
soluble  in  carbon  disulphide,  and  smells  strongly  of  garlic.  Yellow 
arsenic  changes  to  the  metallic  form  at  room  temperature;  the 
change  is^markedly  accelerated  by  exposure  to  light.  Other  modifica- 
tions of  arsenic  have  been  described.1 

The  vapour  density  of  arsenic  at  600  to  700°  corresponds  with  the 
formula  As4,  but  it  gradually  diminishes  with  increasing  temperature, 
and  at  1700°  has  fallen  to  half  the  original  value  ;  the  formula  is 
then  As2. 

When  heated  in  the  air  arsenic  burns  to  the  trioxide,  As2O3.    When 

1  Cf.  Erdmann  and  Reppert,  Annalen,  1908,  361,  i. 


ARSENIC  513 

fragments  of  arsenic  are  thrown  into  chlorine  or  bromine  at  room  tem- 
perature, a  vigorous  reaction  occurs,  and  the  trihalide  (AsCl3 ;  AsBr3) 
is  formed. 


COMPOUNDS  OF  ARSENIC  WITH  HYDROGEN  AND  WITH  THE 
HALOGENS 

Arsine  {Arseniuretted  Hydrogen),  AsH3  —  Preparation  —  (i) 
Arsine  is  obtained,  mixed  with  hydrogen,  when  a  soluble  compound 
of  arsenic  is  added  to  a  solution  in  which  hydrogen  is  being  generated, 
e.g.  by  the  action  of  hydrochloric  or  sulphuric  acid  on  zinc.  It  is  also 
obtained  when  the  hydrogen  is  generated  electrolytically  in  a  solution 
containing  a  compound  of  arsenic  : 


(2)  The  pure  compound  is  prepared  by  the  action  of  dilute  sulphuric 
acid  on  zinc  or  sodium  arsenide  : 

Zn3As2  +  6HCl->3ZnCl2  -H  2AsH3. 

Sodium  arsenide  is  easily  obtained  by  heating  to  redness  a  mixture  of 
metallic  sodium  and  excess  of  arsenious  oxide. 

Properties  —  Arsine  is  a  colourless  gas  with  an  odour  of  garlic, 
and  is  extremely  poisonous.  It  can  be  condensed  to  a  colourless 
liquid,  which  boils  at  -  55°. 

Arsine  is  an  endothermic  compound,  and  is  readily  decomposed 
into  its  elements  by  heat.  This  is  readily  shown  by  passing  the 
mixture  of  arsine  and  hydrogen  obtained  in  (i)  through  a  hard  glass 
tube  and  heating  the  latter  at  a  definite  point,  when  a  mirror  of 
arsenic  will  form  in  the  tube  just  beyond  the  heated  portion. 

Arsine  burns  in  air  with  a  vivid  blue  flame  to  the  trioxide  and 
water  : 


but  when  the  supply  of  air  is  insufficient  it  burns  to  arsenic  and  water: 


The  reactions  just  mentioned  form  the  basis  of  Marsh's  test  for  the 
detection  of  arsenic.  The  apparatus  used  is  shown  in  Fig.  88. 
Hydrogen  is  generated  in  the  Woulfs  bottle  from  zinc  and  dilute 
sulphuric  acid  (which  must  themselves  be  arsenic-free).  When  all 
the  air  has  been  displaced  from  the  apparatus,  the  hydrogen  is  lighted 
33 


5 14     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

at  the  end  of  the  tube,  which  is  drawn  out  to  a  jet,  and  the  arsenic 
solution  added  through  the  funnel.  The  hydrogen  flame  soon 
becomes  livid  blue,  and  when  a  white  porcelain  dish  is  held  in 
the  flame  black  spots  of  metallic  arsenic  are  formed.  Under  the 
same  circumstances  antirflony  gives  similar  spots,  but,  unlike  the 
arsenic  spots,  they  are  insoluble  in  a  fresh  solution  of  bleaching 


FIG.  88. 


powder.     The  decomposition  of  arsine  by  heat,  already  referred  to, 
may  be  demonstrated  by  heating  the  tube  as  shown. 

Arsine  is  a  powerful  reducing  agent  ;  it  readily  precipitates  gold  and 
silver  from  solution  : 


From  a  solution  of  copper  sulphate,  copper  arsenide,  Cu3As2,  is 
precipitated. 

The  Nascent  State—  In  the  previous  section  it  has  been  pointed 
out  that  when  hydrogen  is  generated  in  a  solution  of  arsenious  oxide, 
the  latter  is  reduced  to  arseniuretted  hydrogen.  When,  however, 
gaseous  hydrogen  '(from  a  Kipp's  apparatus,  for  example)  is  passed 
through  the  solution,  no  reduction  of  arsenious  oxide  occurs.  Hydrogen 
generated  in  the  solution  —  the  so-called  "  nascent  "  hydrogen  —  there- 
fore appears  to  be  a  much  more  powerful  reducing  agent  than  ordinary 
hydrogen.  The  usual  explanation  of  this  phenomenon  is  that  the 
hydrogen  at  the  moment  of  liberation  is  in  the  atomic  condition,  and 
is  therefore  much  more  active  chemically. 

The  energy  relations  of  the  systems  throw  some  light  on  this  ques- 


ARSENIC  515 

tion.  In  the  reaction  under  consideration  the  total  heat  change  is 
made  up  of  two  parts  :  (a]  the  heat  of  solution  of  zinc  in  acid,  which 
is  positive  and  large  in  amount  ;  (b)  the  heat  given  out  when  arsenious 
oxide  is  reduced  by  hydrogen,  which  is  negative  although  small. 
The  whole  reaction  is  therefore  exothermic,  whereas  the  reduction 
of  arsenious  oxide  by  free  hydrogen  is  endothermic;  and  it  is  there- 
fore to  be  anticipated  that,  quite  apart  from  any  question  as  to  atomic 
hydrogen,  the  former  reaction  will  proceed  much  more  readily  than  the 
latter  (p.  148).  It  is  highly  improbable,  however,  that  the  energy  of 
one  reaction  can  be  transferred  to  another  entirely  independent  reac- 
tion proceeding  in  the  same  system,  and  we  must  therefore  assume  that 
the  reactions  (a)  and  (£)  are  connected  in  some  way  —  perhaps  through 
the  agency  of  an  intermediate  compound  —  in  order  that  the  energy 
given  out  in  reaction  (a)  may  become  available  for  the  whole  change. 
Apart  from  the  question  of  the  total  amount  of  energy,  the  effect  of 
catalytic  agents  on  the  speed  of  reaction  —  in  the  present  case,  for 
instance,  the  effect  of  the  zinc  on  the  activity  of  the  hydrogen  —  is  also 
of  importance;  but  the  matter  cannot  be  further  discussed  at  the 
present  stage. 

Arsenic  Halides  —  The  following  compounds,  AsF3,  AsCl3,  AsBr3 
and  AsI3  are  definitely  known.  The  only  halide  containing  penta- 
valent  arsenic  which  is  definitely  known  is  the  pentafluoride  AsF6. 

Arsenic  Fluoride,  AsF3,  is  formed  by  direct  combination  of 
its  elements,  but  is  most  readily  obtained  by  distilling  a  mixture  of 
arsenious  oxide,  potassium  fluoride,  and  excess  of  concentrated 
sulphuric  acid  from  a  lead  retort: 

As2O3  +  6H  F->2  AsF3  +  3H2O. 

Properties  —  Arsenic  fluoride  is  a  colourless,  fuming  liquid,  which 
boils  at  63°,  and  is  at  once  decomposed  by  water  into  the  trioxide 
and  hydrofluoric  acid.  On  account  of  its  great  activity  it  is  used  in 
preparing  other  fluorides.  It  is  best  kept  in  platinum  bottles. 

Arsenic  Chloride,  AsCl3,  is  formed  by  heating  arsenic  in  a 
current  of  chlorine,  but  is  most  easily  obtained  by  distilling  a  mixture 
of  arsenious  oxide,  sodium  chloride,  and  concentrated  sulphuric  acid  : 


Properties  —  Arsenic  chloride  is  a  colourless,  fuming  liquid,  which 
boils  at  130°,  and  is  very  poisonous.  It  is  decomposed  reversibly  by 
water,  with  formation  of  the  oxide  and  hydrogen  chloride: 


516     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

With   a   small   quantity   of  water  an   oxychloride,  AsOCl,H2O,   or 
As(OH)2Cl,  is  obtained  as  a  white  precipitate. 

Arsenic  Bromide,  AsBr3,  and  Arsenic  Iodide,  AsI3,  are 
most  easily  obtained  by  adding  powdered  arsenic  to  a  solution  of  the 
halogen  in  carbon  disulphide.  The  bromide  occurs  in  colourless 
prismatic  crystals,  which  melt  at  31°  ;  the  liquid  boils  at  221°.  The 
iodide  forms  red,  lustrous,  hexagonal  crystals,  which  melt  at  146°. 


OXIDES  AND  OXYACIDS  OF  ARSENIC  . 
Two  oxides  of  arsenic  are  known,  both  of  which  are  acidic  : — 

Arsenious  oxide          .        .        .       As2O3  (or  As4O6) 
Arsenic  oxide     ....      As2O6 

No  acid  corresponding  with  arsenious  oxide  has  been  isolated,  but 
salts  derived  from  orthoarsenious  acid,  H3AsO3,  and  from  metarsenious 
acid,  HAsO2,  are  known.  From  arsenic  oxide  three  acids,  of  the 
same  type  as  the  phosphoric  acids,  are  derived  : 

Orthoarsenic  acid H3AsO4 

Pyroarsenic  acid      .....       H4As2O7 
Metarsenic  acid HAsO3 

Arsenious  Oxide,  As2O3 — This  compound,  also  known  as  "  white 
arsenic "  or  "  arsenic,"  is  the  most  important  compound  of  arsenic, 
and  was  familiar  to  chemists  in  the  Middle  Ages.  It  occurs  naturally 
as  arsenolite,  and  is  formed  when  the  element  burns  in  air  or  oxygen. 
It  is  obtained  commercially  in  the  roasting  of  arsenical  pyrites  (iron 
oxide  remaining  behind)  and  of  other  pyrites  containing  arsenic 
(often  as  secondary  product),  the  vapours  being  condensed  in  long 
tubes.  It  is  purified  by  sublimation  from  iron  vessels  connected  with 
cylindrical  receivers,  in  which  it  condenses  in  the  vitreous  form. 

Properties — Arsenious  oxide  exists  in  three  allotropic  modifications, 
two  of  which  are  crystalline  and  one  amorphous. 

(1)  The  octahedral  modification  (regular  system).     This  modifica- 
tion, which   is   stable  at  room  temperature,  is  obtained  by  rapidly 
cooling  the  vapour  of  the  oxide,  or  (in  well-formed  crystals)  by  allow- 
ing a  solution  of  the  oxide  in  hydrochloric  acid  to  crystallize.     The 
density  is  about   3.6.      At  15°  100  grams  of  water   dissolve  about 
1.65  grams,  at  25°  about  2.04  grams  of  this  form,  but  the  rate  of 
solution  is  extremely  slow.' 

(2)  The  prismatic  form  is  obtained  by  allowing  a  solution  of  the 


ARSENIC  517 

oxide  in  potassium  hydroxide  to  evaporate  slowly  in  the  absence  of 
nuclei  of  the  octahedral  form.  This  form  is  metastable  at  room 
temperature;  its  density  is  about  4.15. 

(3)  The  amorphous  modification  is  obtained  as  a  colourless,  trans- 
parent, glassy  mass  by  slow  condensation  of  the  vapour  of  the  oxide 
at  a  relatively  high  temperature.  It  is  unstable  at  room  temperature, 
and  slowly  becomes  opaque  owing  to  its  transformation  to  the 
octahedral  form.  As  the  change  proceeds  from  without  inwards, 
masses  of  oxide  are  often  met  with  which  consist  of  the  octahedral 
form  on  the  outside  with  a  core  of  the  vitreous  form.  In  accordance 
with  the  general  rule  (p.  241),  the  vitreous  form  is  the  more  soluble 
in  water,  but  the  exact  solubility  cannot  be  determined  owing  to 
the  fact  that  the  transformation  is  accelerated  by  water.  The  density 
of  the  vitreous  form  is  about  3.71. 

Arsenious  oxide  passes  directly  into  vapour  on  heating,  but  can  be 
melted  under  pressure.  The  density  of  the  vapour  up  to  1500°  corre- 
sponds with  the  formula  As4O6,  at  1800°  with  the  formula  As2O3. 

The  aqueous  solution  of  arsenious  oxide  is  slightly  acid,  indicating 
the  presence  of  arsenious  acid,  H3AsO3  (or  perhaps  HAsO2). 

Arsenic  trioxide  is  a  powerful  poison.  It  is  remarkable  that  when 
it  is  taken  in  gradually  increasing  doses,  the  human  organism  acquires 
increased  resistance  to  its  action,  so  that  quantities  may  ultimately 
be  taken  without  danger  which  would  at  once  prove  fatal  to  one 
unused  to  it.  It  is  employed  as  a  rat  poison,  in  calicqpprinting, 
etc.,  and  is  also  used  in  medicine. 

Arsenious  Acid  and  Arsenites— Whether  the  acidity  of  a 
solution  of  arsenious  oxide  in  water  is  due  to  the  presence  of 
orthoarsenious  acid,  H3AsO3,  or  of  meta-arsenious  acid,  HAsO2, 
has  not  been  definitely  established.  Arsenious  acid  has  not  been 
obtained  in  the  free  condition  ;  on  evaporating  the  solution  the 
anhydride  separates.  It  is  a  very  weak  acid,  about  the  same 
strength  as  boric  acid  (p.  376). 

The  salts  of  arsenious  acid,  the  arsenites,  are  derived  from  the 
ortho  acid,  H3AsO3,  the  meta  acid,  HAsO2,  the  pyro  acid,  H4As2O5, 
and  still  more  complex  acids.  The  most  important  are  silver 
arsenite,  Ag3AsO3,  which  is  yellow,  and  copper  hydrogen  arsenite 
or  Scheele's  green,  CuHAsO3.  The  latter  is  used  as  a  pigment. 

Arsenic  Pentoxide,  As2Or>,  is  obtained  by  heating  the  tri- 
oxide with  nitric  acid,  and  dehydrating  the  arsenic  acid  thus 
obtained  by  heating  to  low  redness  : 


5i8    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Properties — Arsenic  pentoxide  is  a  white,  amorphous,  deliquescent 
solid,  which  dissolves  in  water  to  form  arsenic  acid.  On  heating 
strongly,  it  decomposes  into  the  trioxide  and  oxygen. 

Arsenic  Acid,  H3AsO4,  and  Arsenates— Arsenic  acid  is 
most  readily  obtained  by  heating  arsenious  acid  with  nitric  acid; 
on  evaporating  the  solution  to  dryness,  and  recrystallizing  the  resi- 
due from  water,  colourless  crystals  of  the  composition  2H3AsO4,H2O 
are  obtained.  At  100°,  the  water  of  crystallization  is  driven  off  and 
the  anhydrous  acid  remains. 

At  140  to  1 80°  two  molecules  of  acid  lose  a  molecule  of  water  and 
pyroarsenic  acid  is  obtained ;  a?  200°  still  more  water  is  expelled 
and  meta-arsenic  acid  results  ;  finally,  at  a  low  red  heat  all  the 
water  is  driven  off  and  the  pentoxide  remains. 

The  analogy  with  the  phosphoric  acids  is  obvious  (cf.  p.  251). 
However,  pyro-  and  meta-arsenic  acids,  unlike  the  corresponding 
phosphorus  compounds,  change  immediately  to  the  ortho  acid  when 
dissolved  in  water,  and,  further,  meta-arsenic  acid,  unlike  meta- 
phosphoric  acid,  can  be  dehydrated  by  heat. 

The  arsenates,  derived  from  the  three  acids  just  mentioned,  closely 
resemble  the  corresponding  phosphates  ;  for  example,  phosphates 
are  in  many  cases  isomorphous  with  the  corresponding  arsenates. 
The  meta-  and  the  pyro-arsenates  change  immediately  to  ortho- 
arsenates  on  dissolving  in  water.  Ammonium  magnesium  arsenate, 
MgNH4A.sO4,  like  the  corresponding  phosphate,  is  insoluble  in 
water,  and  advantage  is  taken  of  this  in  estimating  arsenic  in 
solution.  Silver  arsenate,  Ag3AsO4,  is  chocolate  in  colour,  where- 
as silver  phosphate  is  yellow,  and  this  test  serves  to  distinguish 
phosphates  and  arsenates. 

COMPOUNDS  OF  ARSENIC  AND  SULPHUR 
Three  sulphides  of  arsenic  are  known  : 

Arsenic  disulphide  (realgar)  .  .  .  As2S2 
Arsenic  trisulphide  (orpiment)  .  .  .  As2S3 
Arsenic  pentasulphide  ....  As2S6 

Arsenic  Disulphide,  As2S2,  occurs  naturally  as  realgar,  and 
is  also  obtained  by  fusing  together  arsenic  and  sulphur  in  the 
calculated  proportions.  On  the  large  scale  it  is  obtained  (con- 
taminated with  trioxide)  by  heating  together  iron  pyrites  and 
arsenical  pyrites : 

2FeS2  +  2FeSAs->4FeS  +  As2Sa. 


ARSENIC  5i9 

Properties  —  Arsenic  disulphide  is  a  red  crystalline  solid  of  density 
3.5  ;  it  was  formerly  used  as  a  paint.  A  mixture  of  realgar,  sulphur 
and  potassium  nitrate  is  used  in  pyrotechny  under  the  name  of  Bengal 
fire  or  Greek  white  fire.  It  burns  in  air  to  the  trioxide  and  sulphur 
dioxide. 

Arsenic  Trisulphide,  As2S3,  occurs  naturally  as  orpiment  in 
yellow  monoclinic  crystals.  It  is  obtained  by  fusing  a  mixture  of  its 
components  in  the  calculated  proportions,  and  is  formed  as  a  yellow 
precipitate  when  hydrogen  sulphide  is  passed  through  an  acidified 
solution  of  an  arsenic  compound,  e.g.  As2O3  : 


Properties  —  When  freshly  precipitated  from  solution,  arsenic 
sulphide  is  amorphous,  and  dissolves  to  a  considerable  extent  in 
cold  water,  forming  a  colloidal  solution.  When  hydrogen  sulphide 
is  passed  into  an  aqueous  solution  of  the  trioxide  in  the  absence 
of  acid,  the  solution  becomes  yellow  but  no  precipitation  occurs, 
the  sulphide  remaining  in  colloidal  solution.  It  is  at  once  pre- 
cipitated by  the  adoption  of  acids  or  salts  (cf.  p.  371). 

Arsenic  trisulphide  is  soluble  in  colourless  ammonium  sulphide 
solution,  with  formation  of  ammonium  thioarsenite  : 

As2S3  +  3(NH4)2S-*2(NH4)3AsS3. 

It  is  also  soluble  in  alkali  hydroxide,  a  mixture  of  arsenite  and  thio- 
arsenite being  formed  : 


The  thioarsenites,  which  may  be  regarded  as  analogues  of  the 
arsenites  with  the  oxygen  replaced  by  sulphur,  are  decomposed 
on  addition  of  hydrochloric  acid,  the  trisulphide  being  precipitated  : 

2(NH4)3AsS3+6HCl^As2S3+6NH4Cl  +  3H2S. 

If  instead  of  colourless  ammonium  sulphide  the  yellow  sulphide  is 
used  to  dissolve  arsenic  trisulphide,  the  solution  contains  ammonium 
thioarsenate,  (NH4)3AsS4,  the  analogue  of  ammonium  arsenate  : 

As2S3  +  3(NH4)2S  +  2S->2(NH4)3AsS4. 

On  adding  excess  of  hydrochloric  acid  to  the  solution,  arsenic  penta- 
sulphide  is  precipitated  : 

2(N  H4)3AsS4  +  6H  C1-5-6N  H4CH-  As2S6  +  3H2S. 


520     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Neither  thioarsenious  acid,  H3AsS3,  nor  thioarsenic  acid,  H3AsS4, 
is  known  in  the  free  condition.  It  is  evident  that  the  sulphides 
As2S3  and  As2S6  bear  the  same  relationship  to  these  hypothetical 
acids  as  As2O3  and  As2O5  do  to  arsenious  and  arsenic  acids. 

Arsenic  Pentasulphide,  As2S5,  is  prepared  as  described 
above  or  by  heating  a  mixture  of  the  elements  in  the  theoretical 
proportions.  It  is  a  light  yellow  solid  which  can  be  fused  and  sub- 
limed without  apparent  decomposition.  It  dissolves  in  ammonium 
sulphide  to  form  ammonium  thioarsenate,  and  in  a  solution  of 
potassium  hydroxide  to  form  a  mixture  of  potassium  arsenate  and 
thioarsenate  : 


Tests  for  Arsenic  —  The  more  important  tests  have  been 
mentioned  in  connexion  with  the  different  compounds.  They  will 
be  further  referred  to  in  connexion  with  antimony,  which  resembles 
arsenic  very  closely  in  many  respects. 

ANTIMONY 

Symbol,  Sb.         Atomic  weight,  120.2.         Molecular  weight,  120.2  (at  1800°). 

History  —  The  chief  ore  of  antimony,  stibnite^  Sb2S3,  has  been  in 
use  from  very  early  times  ;  it  is  mentioned  in  the  Old  Testament  as 
being  employed  by  women  for  darkening  their  eyebrows.  A  vessel 
used  by  the  ancient  Chaldeans  has  been  shown  to  be  made  of  metallic 
antimony  (Berthelot).  Stibnite  was  called  stibium  by  the  Romans, 
hence  the  symbol  for  the  element  still  in  use.  In  the  Middle  Ages, 
compounds  of  antimony  were  largely  used  for  medicinal  purposes. 

Occurrence  —  Rich  deposits  of  metallic  antimony  of  a  fairly  high 
degree  of  purity  have  recently  been  found  in  Australia.  In  combina- 
tion with  oxygen,  as  Sb2O3,  it  constitutes  the  mineral  senarmontite, 
and  as  Sb2O4  it  constitutes  antimony  ochre.  The  most  important  ore, 
however,  and  the  one  from  which  the  metal  is  chiefly  obtained,  is 
stibnite  or  gray  antimony  ore,  Sb2S3.  A  large  number  of  ores  contain- 
ing antimony,  sulphur,  and  either  lead,  copper,  silver  or  iron  are 
also  met  with. 

Preparation  Of  Metal  —  The  first  step  in  preparing  antimony 
from  stibnite  is  to  heat  the  ore  in  an  earthenware  pot,  when  the  sulphide 
fuses  and  is  drained  away  (through  holes  in  the  bottom  of  the  pot) 
from  the  accompanying  impurities.  From  this  partially  purified 
sulphide  the  metal  is  prepared  by  one  of  two  processes. 


ANTIMONY  521 

(1)  The  sulphide  is  heated  with  scrap  iron,  whereby  ferrous  sulphide 
is  formed  and  floats  on  the  top,  the  antimony  settling  out  below. 

Sb2S3  4-  3Fe->3FeS  +  2Sb. 

(2)  According  to  the  second  method  the  ore,  mixed  with  charcoal 
to  prevent  it  from  caking,  is  roasted  to  convert  it  to  the  oxide  (the 
tetroxide,  Sb2O4,  is  the  chief  product)  ;  it  is  then  mixed  with  a  further 
quantity   of  charcoal   and   strongly   heated,   whereby  the   metal   is 
liberated  : 

Sb2S3  +  5O2->Sb2O4  +  38  O2 


The  antimony  thus  obtained  contains  copper,  iron,  arsenic  and  other 
impurities.  It  can  be  purified  by  fusing  with  a  mixture  of  antimony 
sulphide  and  sodium  carbonate.  The  sulphides  of  the  other  metals, 
formed  by  double  decomposition,  rise  to  the  surface  and  can  be 
removed. 

Properties  —  Antimony  exists  in  at  least  three  different  allotropic 
modifications  ;  (a)  metallic  antimony,  the  ordinary  stable  form,  which 
occurs  in  hexagonal  crystals  ;  (b)  explosive  antimony  ;  (c]  yellow 
antimony. 

(a)  Metallic  antimony  (the  ordinary  form)  is  a  silver-white,  lustrous, 
brittle  metal  of  density  6.5  ;  it  melts  at  630.6°,  and  boils  about  1300°. 
It  conducts  heat  and  electricity,  but  is  a  much  less  efficient  conductor 
than  the  typical  metals. 

(b}  Explosive  antimony  is  obtained  as  a  black  shining  substance 
on  the  cathode  when  a  solution  of  antimony  trichloride  in  hydrochloric 
acid  is  electrolyzed  with  an  antimony  anode  and  platinum  cathode. 
Its  density  is  5.78.  It  is  unstable  under  ordinary  conditions,  and  ex- 
plodes vigorously  when  scratched  with  a  needle  or  rubbed  in  a  mortar, 
with  formation  of  ordinary  antimony.  It  always  contains  a  greater  or 
less  amount  of  adsorbed  antimony  trichloride.  Whether  explosive 
antimony  is  crystalline  or  amorphous  has  not  been  established. 

(c]  Yellow  antimony  is  obtained  by  the  action  of  oxygen  on  liquid 
antimony  hydride,  SbH3,  at  —  90°  : 


Like  the  corresponding  modifications  of  phosphorus  and  arsenic,  it  is 
soluble  in  carbon  disulphide.  It  is  very  unstable,  and  slowly  changes 
to  the  metallic  modification  even  at  -90°. 

The  vapour  density  of  antimony  at  1400°  exceeds  that  required  by 


522     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

the  formula  Sb2,  but  at  1800°  it  is  about  60,  indicating  that  the  metal 
under  these  conditions  is  monatomic. 

Antimony  is  stable  in  the  air,  but  on  heating  burns  to  the  trioxide. 
It  combines  directly  with  chlorine  and  bromine  at  room  temperature. 
It  is  slowly  dissolved  on  boiling  with  concentrated  hydrochloric  acid, 
with  evolution  of  hydrogen.  When  heated  with  concentrated  sul- 
phuric acid,  antimony  sulphate,  Sb2(SO4)3,  is  formed  and  sulphur 
dioxide  given  off.  Dilute  nitric  acid  oxidizes  antimony  chiefly  to 
the  trioxide,  Sb2O3;  with  concentrated  nitric  acid  the  pentoxide, 
Sb2O6  (or  one  of  the  corresponding  acids),  is  the  chief  product. 

Alloys — The  most  important  alloys  of  antimony  are  type  metal 
(which  contains  antimony,  tin  and  lead  in  varying  proportions)  and 
Britannia  metal  (which  contains  tin,  antimony  and  copper).  Antimony 
expands  on  solidification,  and  therefore  the  alloys  containing  it,  when 
used  for  casting,  give  sharp  impressions. 

Antimony  Hydride  (Stibine)  SbH3 — The  methods  of  prepara- 
tion and  properties  of  this  compound  are  very  similar  to  those  of 
arsine  (p.  513).  It  is  formed  (i)  when  a  soluble  antimony  compound 
is  added  to  a  mixture  of  zinc  and  dilute  sulphuric  acid  ;  (2)  in 
much  purer  condition  by  the  action  of  dilute  hydrochloric  acid  on 
magnesium  antimonide. 

Properties — Antimony  hydride  is  a  colourless  gas  with  a  character- 
istic musty  odour.  The  liquefied  gas  boils  at  -  18°,  the  solid  melts  at 
-91°.  It  is  an  endothermic  compound,  and  slowly  decomposes  into 
its  elements  even  at  room  temperature  ;  the  decomposition  is  greatly 
accelerated  by  a  mirror  of  antimony  deposited  on  the  tube.  It  burns 
in  air  with  a  livid-blue  flame  to  the  trioxide  and  water,  but  when  the 
supply  of  air  is  insufficient,  water  and  antimony  are  the  first  products. 
When  therefore  a  porcelain  dish  is  held  in  the  flame,  a  black  spot  of 
antimony  is  obtained  which,  unlike  the  corresponding  arsenic  spot 
(p.  514),  is  insoluble  in  bleaching  powder  solution.  When  passed  into 
a  solution  of  silver  nitrate,  silver  antimonide,  SbAg3,  is  precipitated  : 

SbH3+3AgN03-»3HN03  +  SbAg3. 

Compounds   of  Antimony   •with  the  Halogens — The 

following  halides  of  antimony  are  known  : 

SbF3  SbCl3  SbBr3  SbI3 

SbF6  SbCl5 

Antimony  Trifluoride,  SbF3,  is  obtained  by  dissolving  the 
trioxide  in  hydrofluoric  acid,  and  separates  in  colourless,  transparent, 


ANTIMONY  523 

rhombic  crystals  when  the  solution  is  evaporated.  It  does  not  fume 
in  moist  air,  and,  unlike  the  other  trihalides,  forms  a  clear  solution 
with  water. 

Antimony  Pentafluoride,  SbF5,  is  best  prepared  by  boiling 
antimony  pentachloride  with  anhydrous  hydrofluoric  acid  for  some 
days  ;  the  residue  is  then  subjected  to  fractional  distillation  : 

SbCl6+  5HF->SbF6+  sHCl. 

The  pentachloride  is  a  colourless,  thick,  oily  iliquid  which  boils  at 
149-150°,  and  forms  a  clear  solution  with  water. 

Both  the  trifluoride  and  pentafluoride  form  complex  compounds 
with  the  alkali  fluorides,  e.g.  SbF3,KF  ;  SbF3,2KF  ;  SbF6,KF  ; 
SbF5,2KF,2H2O. 

Antimony  Trichloride,  SbCl3,  is  obtained  when  chlorine  is 
passed  over  metallic  antimony,  but  is  most  readily  obtained  by  dis- 
solving antimony  oxide  or  sulphide  in  concentrated  hydrochloric 
acid  : 

Sb2S3+6HCl->2SbCl3  +  3H2S. 

Properties  —  Antimony  trichloride  forms  a  soft  crystalline  mass, 
which,  in  allusion  to  its  consistency,  is  known  as  "butter  of  antimony." 
It  melts  at  73.2°  to  an  oily  liquid,  which  boils  at  223°. 

With  a  small  quantity  of  water  a  clear  solution  is  obtained,  but 
when  more  water  is  added  hydrolysis  occurs,  and  one  or  both  of  the 
oxychlorides  SbOCl  and  (SbOCl)2,Sb2O3  are  obtained,  depending 
upon  the  conditions  : 


4SbCl3+5H2O^(SbOCl)2,Sb2O3+ioHCl. 

The  first  is  the  main  reaction  in  cold,  the  second  in  hot  solutions. 
By  boiling  with  successive  quantities  of  water  complete  hydrolysis 
can  be  effected  : 

(SbOCl)2,Sb2O3+H2O->2Sb2O3+2HCl. 

Antimony  Pentachloride,  SbCl5,  is  obtained  when  antimony 
burns  in  excess  of  chlorine,  but  is  best  prepared  by  passing  chlorine 
into  the  trichloride  : 

SbCl3  +  Cl2-»SbCl6. 

Properties  —  Antimony  pentachloride  is  a  colourless  (as  commonly 
met  with  often  yellowish)  fuming  liquid,  which  can  be  obtained  in 


524    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

crystals  melting  at  -  6°.     When  heated,  it  boils  about  140°,  the  vapour 
being  partly  dissociated  into  the  trichloride  and  chlorine : 


Under  reduced  pressure  it  can  be  distilled  without  decomposition. 

With  small  quantities  of  water  it  forms  two  crystalline  hydrates, 
SbCl5,H2O  and  SbCl6,4H2O.  When  more  water  is  added,  it  undergoes 
hydrolysis,  oxychlorides,  probably  SbOCl3  and  SbO2Cl,  being  precipi- 
tated. With  excess  of  hot  water  it  is  completely  decomposed,  with 
formation  of  hydrochloric  and  antimonic  acids  (cf.  PC16,  p.  1246) : 

SbClfi  +  4H2O->SbO(OH)3+5HCl. 

Antimony  pentachloride  forms  a  large  number  of  crystalline  "  addi- 
tion compounds"  with  the  chlorides  of  other  elements,  and  with 
organic  compounds,  e.g.  SbCl5,KCl,H2O  ;  SbCl6,NH4Cl,H2O  ; 
SbCl5,FeCl3,8H2O  ;  SbClg^HCN. 

Antimony  Tribromide,  SbBr3,  and  the  Tri-iodide,  SbI3, 
are  prepared,  like  the  corresponding  arsenic  compounds,  by  adding 
powdered  antimony  to  bromine  or  iodine  dissolved  in  carbon  disul- 
phide.  The  bromide  forms  colourless,  rhombic  crystals,  which  melt 
at  90  to  94° ;  the  liquid  boils  at  275°.  The  iodide  occurs  in  three 
allotropia  modifications,  the  most  stable  forming  ruby-red  crystals. 
Both  halides  are  hydrolyzed  by  water. 

OXIDES  AND  OXYACIDS  OF  ANTIMONY 
Three  oxides  of  antimony  are  known : 

Antimony  trioxide  (antimonious  oxide)     .     Sb2O3  (or  Sb4O6). 

Antimony  tetroxide Sb2O4. 

Antimony  pentoxide          ....     Sb2O6. 

From  the  trioxide  is  derived  Sb(OH)3,  or  H3SbO3,  which  has  both 
acidic  and  basic  properties.  The  acidic  properties  are  shown  in  the 
existence  of  the  compound,  NaSbO2,3H2O  (octahedral  crystals)  which, 
however,  is  derived  from  metantimonious  acid,  HSbO2.  The  basic 
character  of  the  hydroxide  is  referred  to  below. 

From  the  pentoxide  three  acids  are  derived  : 

'  Orthoantimonic  acid        .         .     H3SbO4. 
Pyroantimonic  acid  .        .     H4Sb2O7. 

Metantimonic  acid  .         .         .     HSbO8. 


ANTIMONY  525 

Antimony  Trioxide,  Sb2O3,  is  obtained,  mixed  with  the 
tetroxide,  when  antimony  is  burned  in  the  air  ;  it  is  best  obtained 
by  boiling  the  trichloride  with  a  dilute  solution  of  sodium  carbonate  : 


Properties  —  Antimony  trioxide  is  a  colourless  substance  which 
exists  in  two  crystalline  modifications  —  (a)  a  rhombic  form,  density 
5.6  ;  (b]  an  octahedral  form,  density  5.3  (cf.  arsenic  trioxide).  It  is 
practically  insoluble  in  water,  and  in  dilute  nitric  or  sulphuric  acid, 
but  dissolves  in  hydrochloric  acid  to  form  the  trichloride,  and  also 
in  a  solution  of  potassium  hydrogen  tartrate,  KHC4H4OC,  to  form 
potassium  antimony  tartrate  (tartar  emetic)  : 

Sb2O3+2KHC4H4O6->2K(SbO)C4H4O6+H2O. 

It  dissolves  in  a  concentrated  solution  of  alkali  hydroxide  to  form  an 
antimonite  (see  above). 

The  univalent  SbO  group,  which  is  met  with  in  antimony 
oxy  chloride,  SbOCl,  and  in  tartar  emetic,  is  further  referred  to 
under  antimonious  hydroxide  (g.v.}. 

Antimony  Tetroxide,  Sb2O4,  is  obtained  by  heating  the  metal, 
the  trioxide,  or  the  pentoxide  in  air  at  a  high  temperature. 

Properties  —  Antimony  tetroxide  is  a  white  powder,  insoluble  in 
water.  It  cannot  be  fused  or  volatilized;  when  heated  to  1000°  it 
decomposes  into  the  trioxide  and  oxygen.  In  many  respects  it 
behaves  like  a  mixture  of  the  trioxide  and  pentoxide,  and  may  be 
regarded  as  antimonious  antimonate,  Sb2O3*Sb2O6. 

Antimony  Pentoxide,  Sb2O6,  is  obtained  by  the  action  ot 
nitric  acid  on  metallic  antimony,  and  heating  the  antimonic  acid  thus 
obtained  to  a  temperature  not  exceeding  300°. 

Properties  —  Antimony  pentoxide  is  a  yellow,  infusible  powder, 
insoluble  in  water.  When  heated  to  440°  it  decomposes  into  the 
tetroxide  and  oxygen.  It  has  weakly  acidic  properties,  but  is  devoid 
of  basic  properties. 

Antimonic  Acids  —  Three  antimonic  acids,  orthoantimonic 
acid,  H3SbO4,  pyroantimonic  acid,  H4Sb2O7,  and  metantimonic  acid, 
HSbO3,  are  known,  corresponding  with  the  phosphoric  and  arsenic 
acids.  The  ortho  acid  is  obtained  by  acting  on  antimony  with  nitric 
acid,  and  heating  the  product  at  100°.  When  orthoantimonic  acid  is 
heated  at  200°  for  some  time,  it  loses  a  molecule  of  water,  and  the 
pyro  acid  is  formed  ;  at  a  slightly  higher  temperature  more  water  is 


526     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

driven  off  and  the  meta  acid,  HSbO3,  results.  The  three  acids  are 
white  powders,  and  show  no  differences  in  chemical  behaviour. 

The  antimonates  are  mainly  derived  from  metantimonic  acid, 
HSbO3,  but  some  salts  of  pyroantimonic  acid  are  also  known. 
When  antimony  is  fused  with  potassium  nitrate,  potassium  met- 
antimonate,  KSbO3,  is  formed  as  a  white  powder.  When  this  salt 
is  fused  with  potassium  hydroxide,  potassium  pyroantimonate  is 
formed : 

2KSbO3 + 2KOH-»K4Sb2O7  +  H2O. 

When  boiled  with  water,  potassium  pyroantimonate  is  decomposed 
into  free  alkali  and  dipotassium  pyroantimonate,  K2H2Sb2O7 : 

K4Sb2O7  +  2  H2O->K2H2Sb2O7  +  2KOH. 

The  latter  salt  is  fairly  soluble  in  water  and  is  used  as  a  test  for 
sodium  ;  its  use  for  this  purpose  is  based  upon  the  fact  that  the 
corresponding  sodium  salt,  Na2H2Sb2O7,6H2O,  is  only  slightly 
soluble  in  water. 

Antimonious  Hydroxide,  Sb(OH)3,  and  its  Salts— Refer- 
ence has  already  been  made  to  the  acidic  properties  of  the  hydroxide, 
Sb(OH)3.  Its  basic  character  is  shown  by  the  existence  of  salts 
derived  either  from  the  acid  itself,  e.g.  Sb(NO3)3;  Sb2(SO4)3,  or  from 
its  first  anhydride,1  SbO(OH),  e.g.  SbOCl. 

Antimony  sulphate,  Sb2(SO4)3,  is  obtained  by  dissolving  the  trioxide 
in  concentrated  sulphuric  acid,  and  separates  from  solution  in  long, 
colourless,  lustrous  needles.  It  is  decomposed  by  water  with  formation 
of  basic  salts.  The  nitrate,  Sb(NO3)3,  obtained  by  dissolving  the 
trioxide  in  cold,  fuming  nitric  acid,  also  occurs  in  colourless  crystals. 
On  heating  gently,  antimony  pentoxide  is  formed.  It  is  hydrolyzed 
by  water.  The  existence  of  these  two  salts  is  conclusive  proof  of 
the  basic  character  of  antimony  in  its  trivalent  compounds. 

Antimony  Sulphides — Two  sulphides  of  antimony  are  known : 
the  trisulpkide,  Sb2S3,  and  the  pentasulphide,  Sb2S6.  Antimony 
trisulphide,  Sb2S3,  is  obtained  as  a  grayish-black  mass  by  fusing 
together  sulphur  and  antimony  in  absence  of  air,  and  as  an  orange- 
red  precipitate  on  passing  hydrogen  sulphide  into  a  solution  of  an 
antimony  salt.  It  dissolves  in  concentrated  hydrochloric  acid  to  form 
the  trichloride,  is  soluble  in  colourless  ammonium  sulphide  solution 
to  form  ammonium  thioantimonite,  (NH4)3SbS3,  and  dissolves  in 

1  The  univalent  group-Sb— O  is  sometimes  termed  the  antimonyl group. 


BISMUTH  527 

yellow  ammonium   sulphide   solution   to  form  ammonium  thioanti- 
monate, (NH4)3SbS4  (cf.  arsenic  sulphides,  p.  519). 

Antimony  Pentasulphide,  Sb2S6,  is  obtained  on  adding  excess 
of  an  acid  to  a  solution  of  a  thioantimonate,  e.g.  (NH4)3SbS4: 


It  is  a  dark  orange-red,  amorphous  powder  ;  on  heating  to  200°  it 
decomposes  into  the  trisulphide  and  sulphur. 

Thioantimonious  acid,  H3SbS3,  and  thioantimonic  acid,  H3SbS4, 
are  not  known  in  the  free  condition,  but  a  number  of  thioanti- 
monites  and  thioantimonates  are  known.  Sodium  thioantimonate, 
Na3SbS4,9H2O,  which  forms  colourless  tetrahedral  crystals,  is  known 
as  Schlippe's  salt. 

Tests  for  Antimony  —  As  already  mentioned,  there  is  a  resem- 
blance between  the  behaviour  of  arsenic  and  of  antimony  in  their 
compounds.  Both  give  Marsh's  test,  but  the  arsenic  spot  is  soluble, 
the  antimony  spot  insoluble  in  solution  of  bleaching  powder.  Anti- 
mony trisulphide  is  soluble,  arsenic  trisulphide  practically  insoluble 
in  concentrated  hydrochloric  acid.  When  passed  into  solution  of 
silver  nitrate,  both  arsenic  and  antimony  hydride  produce  black 
precipitates  (p.  522),  but  whereas  the  arsenic  remains  in  the  solution 
as  arsenious  acid,  the  antimony  is  precipitated  as  silver  antimonide, 
SbAg3. 

BISMUTH 

Symbol,  Bi.     Atomic  weight  =208.0.     Molecular  weight  =  208.0. 

Occurrence  —  Bismuth  is  not  a  very  abundant  element,  but  is 
fairly  widely  distributed  in  nature.  It  occurs  chiefly  in  the  free 
condition,  along  with  granite  and  with  cobalt  and  silver  ores.  In 
the  combined  state  it  is  found  as  the  sulphide,  Bi2S3,  bismuthite  or 
bismuth  glance,  as  the  oxide,  Bi2O3,  bismuth  ochre^  and  in  many 
other  forms. 

Preparation  —  Formerly  the  metal  was  obtained  by  heating  the 
ore  containing  it  in  inclined  pipes  when,  owing  to  its  low  melting- 
point,  the  metal  drained  away,  leaving  the  impurities  behind.  Another 
method  now  largely  used  is  to  roast  the  ore  and  heat  with  coal  and  a 
flux.  The  impurities  rise  to  the  surface,  and  the  fused  bismuth  can  be 
separated. 

In  order  to  purify  the  metal  thus  obtained,  it  is  fused  with  potassium 
nitrate  and  sodium  chloride,  whereby  the  impurities  are  oxidized.  A 
more  thorough  purification  is  effected  by  dissolving  the  pure  metal 


528    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

in  nitric  acid  to  form  the  nitrate,  Bi(NO3)3,  and  then  adding  water 
to  the  solution,  whereby  the  oxynitrate,  BiONO3,  is  precipitated.  The 
oxynitrate  is  then  heated  to  convert  it  to  the  oxide,  and  the  latter 
reduced  with  carbon. 

Properties — Bismuth  is  a  white  lustrous  metal  with  a  reddish 
tinge.  Its  density  is  9.78  at  20°;  it  melts  at  about  271°  and  boils 
at  1420°.  It  is  brittle  and  a  rather  inferior  conductor  of  heat  and 
electricity.  It  crystallizes  in  rhombohedral  crystals  belonging  to  the 
hexagonal  system. 

Bismuth  is  stable  in  the  air,  but  on  heating  burns  with  a  bluish 
flame,  forming  the  trioxide.  It  is  scarcely  affected  by  hydrochloric 
acid  unless  oxygen  is  present,  but  dissolves  in  hot  sulphuric  acid  ;  in 
nitric  acid  and  in  aqua  regia  it  dissolves  readily  at  room  temperature. 
In  each  case  compounds  of  trivalent  bismuth  are  formed. 

Bismuth  is  used  as  a  constituent  of  alloys,  many  of  which  fuse  at 
very  low  temperatures.  Rose's  metal  contains  bismuth  2  parts,  lead 
i  part  and  tin  I  part,  and  melts  at  94° ;  Wood's  metal,  which  consists 
of  bismuth  4  parts,  lead  2  parts,  tin  I  part,  cadmium  i  part,  melts  at 
61°.  Bismuth,  like  antimony,  expands  on  solidification,  and  therefore 
the  alloys  containing  it  give  sharp  castings. 


COMPOUNDS  OF  BISMUTH  WITH  THE  HALOGENS 

Only  the  following  compounds,  containing  trivalent  bismuth,  are 
definitely  known : 

BiF3  BiCl3  BiBr3  BiI3. 

Compounds  of  the  type  BiX  and  B5X2  (for  example  BiCl  and  BiCl2) 
have  also  been  described,  but  the  evidence  regarding  their  existence 
is  contradictory. 

Bismuth  Trifluoride,  BiF3,  is  obtained  by  dissolving  bismuth 
trioxide  in  hydrofluoric  acid.  On  evaporating  the  solution,  it  is 
obtained  as  a  grayish-white  crystalline  powder,  which  is  practically 
insoluble  in,  and  is  not  attacked  by,  water. 

Bismuth  Trichloride,  BiCl3,  is  obtained  in  the  anhydrous 
condition  by  passing  dry  chlorine  over  bismuth  gently  heated  in  a 
retort  (cf.  PC13,  p.  245)  ;  it  is  obtained  in  solution  by  dissolving  the 
metal  in  aqua  regia  or  the  trioxide  in  hydrochloric  acid. 

Properties — Bismuth  trichloride  is  generally  met  with  in  colourless 
deliquescent  crystals  which  melt  at  225-230° ;  the  liquid  boils  at 


BISMUTH  529 

447°.     It  is  hydrolyzed  by  water,  bismuth  oxychloride,  BiOCl,  being 
precipitated  : 

BiCl 


Bismuth  chloride  forms  a  large  number  of  complex  compounds  with 
alkali  halides,  e.g.  NaBiCl4,3H2O  or  BiCl3,NaCl,3H2O  ;  K2BiCl6,2H2O 
and  (NH4)2BiCl5,2H2O. 

Bismuth  Tribromide,  BiBr3,  is  prepared  by  slowly  adding 
powdered  bismuth  to  bromine,  allowing  to  stand  for  several  days  and 
then  distilling.  It  forms  yellow,  lustrous  crystals  and  is  decomposed 
by  water  with  formation  of  the  oxybromide  BiOBr. 

Bismuth  Triiodide,  BiI3,  can  be  obtained  by  heating  the 
elements  together  in  an  atmosphere  of  carbon  dioxide,  but  is  best 
prepared  by  heating  a  mixture  of  bismuth  trisulphide  and  iodine  and 
then  subliming  the  triiodide.  The  compound  prepared  by  sublimation 
forms  lustrous,  black  leaflets,  and  is  decomposed  by  water  with  forma- 
tion of  oxyiodide,  BiOI. 

OXIDES  OF  BISMUTH 

Only  the  oxides  BiO  and  Bi2O3  are  known  with  certainty.  Mixtures 
containing  a  higher  proportion  of  oxygen  can  also  be  obtained,  but 
no  definite  compound  has  been  isolated. 

Bismuthous  Oxide,  BiO,  is  obtained  by  heating  the  basic 
oxalate,  (BiO)2C2O4,  in  a  current  of  carbon  dioxide: 

(BiO)2C2O4-»2BiO  +  2CO2. 

Properties  —  Bismuthous  oxide  is  a  black  powder,  which  when 
heated  in  air  combines  with  oxygen  to  form  the  trioxide.  When 
heated  to  350°  in  a  current  of  carbon  dioxide  it  decomposes  into 
bismuth  trioxide  and  metallic  bismuth.  It  acts  as  a  reducing  agent. 

Bismuth  Trioxide,  Bi2O3,  is  formed  when  the  metal  is  burned 
in  aiis  but  is  most  conveniently  prepared  by  heating  the  hydroxide, 
the  basic  nitrate  or  carbonate  : 

2Bi(OH)3-»Bi2O3+  3H2O. 
2BiONO3->Bi2O3+2NO2  +  O. 
(BiO)2CO3-»Bi2O,  +  CO2. 

Properties  —  Bismuth  trioxide  occurs  as  a  yellow,  amorphous  powder, 
or  in  yellow,  lustrous  crystals.  It  is  stable  in  the  air.  It  dissolves  in 
excess  of  acids  to  form  salts  which  contain  Bi'"  ions  : 


34 


530     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Bismuth  Hydroxide,  Bi(OH)3,  is  obtained  as  a  white,  amor- 
phous precipitate  by  pouring  a  solution  of  bismuth  nitrate  into  excess 
of  ammonium  hydroxide: 

Bi(N03)3  +  3NK4OH->Bi(OH)3  +  3NH4N03. 

When  dried  at  100°,  it  loses  water  and  forms  the  first  anhydride, 
BiO(OH)  or  Bi2O3,H2O. 

Bismuth  hydroxide  has  no  acidic  properties,  but  it  acts  as  a  base, 
combining  with  acids  to  form  salts.  The  latter  are  derived  either 
from  the  normal  hydroxide,  B5(OH)3,  or  from  the  compound  BiOOH 
(see  below). 

Higher  Oxides  of  Bismuth  (Peroxides  of  Bismuth}— When 
bismuth  trioxide  is  suspended  in  alkali  and  chlorine  is  passed  into 
the  mixture,  oxides  of  bismuth  are  obtained  which  contain  a  higher 
proportion  of  oxygen  than  the  trioxide.  The  product  containing 
most  oxygen  is  obtained  by  using  a  concentrated  solution  of  potassium 
hydroxide.  After  saturating  with  chlorine,  excess  of  concentrated 
nitric  acid  is  added.  The  insoluble  product  is  yellowish-red  to 
scarlet-red  in  colour,  and  may  be  a  mixture  of  Bi2O4  and  Bi2O6,  but 
no  definite  evidence  on  the  matter  has  so  far  been  obtained.  It  has 
further  been  stated  that  the  production  in  question  has  acidic  properties, 
and  it  has  therefore  been  termed  bismuthic  acid  (from  analogy  with 
antimonic  acid).  On  this  point  also,  however,  the  evidence  is  at 
present  contradictory. 

Normal  and  Basic  Salts — As  mentioned  above,  bismuth  salts 
are  derived  exclusively  from  the  hydroxide,  Bi(OH)3,  and  its  an- 
hydride, BiO(OH),  acting  as  bases.  The  trichloride,  BiCl3,  and 
the  oxychloride,  BiOCl,  have  already  been  discussed.  Bismuth. tri- 
nitrate,  Bi(NO3)3,  is  obtained  by  dissolving  the  metal  in  dilute  nitric 
acid,  and  separates  on  concentrating  the  solution  in  large  triclinic, 
colourless  crystals,  with  5H2O.  The  salt  is  decomposed  by  water 
with  formation  of  the  basic  nitrate,  BiO'NO3,  a  white  powder,  which 
is  used  in  medicine  under  the  name  subnitrate  of  bismuth.  A  number 
of  basic  nitrates  are  known. 

Bismuth  Sulphate,  Bi2(SO4)3,  is  obtained  by  dissolving  the 
trioxide  in  moderately  concentrated  sulphuric  acid,  and  separates 
from  solution  in  colourless,  hygroscopic  needles,  which  are  decom- 
posed by  water  with  formation  of  basic  salts. 

Bismuth  Trisulphide,  Bi2S3,  is  the  only  sulphide  of  bismuth 
definitely  known.  It  is  obtained  in  crystalline  form  (grayish  leaflets) 
by  heating  bismuth  with  excess  of  sulphur,  and  as  a  brownish-black 


ELEMENTS    OF   THE    NITROGEN    GROUP      531 


amorphous  precipitate  by  passing  hydrogen  sulphide  into  a  solution 
of  a  bismuth  salt.  Under  great  pressure  the  amorphous  changes  to 
the  crystalline  modification  ;  increase  of  temperature  greatly  accele- 
rates the  change.  The  trisulphide  is  practically  insoluble  in  solutions 
of  alkali  hydroxides  and  sulphides  (cf.  As2S3  ;  Sb2S3),  but  dissolves  in 
hot  hydrochloric  or  nitric  acid. 

Tests  for  Bismuth — A  valuable  test  for  bismuth  compounds 
is  the  formation  of  a  white  precipitate  of  basic  salt  when  a  solution 
(preferably  of  the  chloride  or  nitrate)  is  poured  into  excess  of  water. 
The  dark  colour  of  the  trisulphide,  and  its  insolubility  in  alkali 
sulphides,  at  once  distinguish  bismuth  compounds  from  those  of 
arsenic  and  antimony. 

General  Characters  of  the  Elements  of  the  Fifth 
Group  and  Summary — The  more  important  physical  properties 
of  the  members  of  the  nitrogen  sub-group  are  shown  in  the  accom- 
panying table,  which  also  illustrates  the  variation  of  these  properties 
with  increasing  atomic  weight : — 


Nitrogen. 

Phosphorus. 

Arsenic. 

Antimony. 

Bismuth. 

Atomic  weight 
Density 
Melting-point 
Boiling-point 

14.01 
i.  026  at  -252° 

-210° 

-195.6° 

31.0 
1.82  to  2.15 

290° 

74.96 
4.72  to  5.73 
ca.  500° 

120.2 

5.78  to  6.5 
360.6° 
ca,  1300° 

208.0 
9.78 
264° 
1420° 

It  will  be  observed  that  bismuth  is  out  of  accord  with  the  remaining 
members  of  the  series  as  regards  its  melting-point  and  boiling-point. 
The  wide  variation  in  the  numbers  quoted  for  the  densities  is  due 
to  the  fact  that  some  of  the  elements  exist  in  different  allotropic 
modifications. 

As  regards  the  chemical  characters,  the  first  point  to  note  is  that 
the  main  valencies  in  the  group  are  three  and  five.  In  the  case  of 
bismuth,  however,  the  existence  of  quinquevalent  compounds  has  not 
been  definitely  proved.  The  valencies  are  not  exclusively  three  and 
five,  however  ;  nitrogen,  in  particular,  functions  with  other  valencies 
as  well.  The  second  point  to  note  is  that  the  non-metallic  character 
gradually  weakens  with  increase  of  atomic  weight;  whilst  nitrogen  is 
a  typical  non-metal,  antimony  and  bismuth  have  distinct  metallic 
characters.  This  change  is  illustrated  by  the  appearance  of  the 
elements  themselves,  as  also  by  the  behaviour  of  their  compounds. 
For  instance,  the  oxides  of  nitrogen  are  exclusively  acidic,  antimony 


532     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

trioxide  is  both  basic  and  acidic,  bismuth  trioxide  is  exclusively  basic. 
The  chlorides  of  nitrogen -and  phosphorus  are  hydrolyzed  irreversibly 
by  water ;  those  of  the  remaining  three  elements  are  hydrolyzed  par- 
tially and  reversibly.  Further,  the  first  four  elements  form  volatile 
hydrates  of  tfie  type  EH3,  which  is  an  indication  of  non-metallic 
character ;  but  they  diminish  in  stability  as  the  atomic  weight  in- 
creases, and  so  far  no  hydride  of  bismuth  has  been  obtained.  The 
existence  of  stable  nitrates  and  sulphates  of  antimony  and  of  bismuth 
is  further  proof  of  basic  properties.  Finally,  the  acidic  properties  of 
the  trisulphides  of  arsenic  and  antimony  is  shown  by  the  existence  of 
thioantimonites,(NH4)3As(Sb)S3  ;  but  bismuth  trisulphide  is  insoluble 
in  alkali  sulphide  solution. 


CHAPTER  XXXIII 

ELEMENTS   OF   THE   CHROMIUM    GROUP   (GROUP  VI, 
SUB-GROUP   A) 

Sub-group  A  Sub-$roup  B 

Chromium,  Cr   .         .        .  .     52.0  Oxygen,  O  16.00 

Molybdenum,  Mo      .         .  .     96.0  Sulphur,  S        .         .        .  .32.07 

Tungsten,  W              .        .  .184.0  Selenium,  Se     .         .        .  .79.2 

Uranium,  Ur     .         .         .  .  238.5  Tellurium          .         .         .  "     .  127.5 

THE  elements  of  sub-group  B  have  already  been  considered  in 
detail  (pp.  20, 289) ;  of  the  four  members  of  sub-group  A  chromium 
is  the  only  important  element.  All  the  members  of  the  sub-group 
exhibit  a  number  of  valencies  ;  the  lower  oxides  are  basic,  the  higher 
acidic.  Corresponding  with  the  position  of  the  elements  in  the  sixth 
group  of  the  periodic  table,  the  typical  oxides  have  the  respective 
formulas  CrO3,MoO3,WO3,UrO3  ;  they  are  almost  exclusively  acidic. 


CHROMIUM 

Symbol,  Cr.     Atomic  weight— 52.0. 

Occurrence — Chromium  does  not  occur  free  in  nature.  The  ex- 
clusive source  of  chromium  compounds  is  chrome  ironstone  or  chromite, 
FeO,Cr2O3 ;  another  natural  compound  of  chromium  is  crocoisite, 
PbCrO4. 

Preparation — The  element  is  now  prepared  commercially 
according  to  the  Goldschmidt  process  (p.  482)  by  reducing  chromic 
oxide  by  means  of  powdered  aluminium,  the  mixture  being  ignited 
by  a  fuse  of  magnesium  ribbon.  When  a  slight  excess  of  the  oxide 
is  used  the  resulting  metal  is  free  from  aluminium. 

Properties — Chromium  is  a  very  hard,  steel-gray  metal,  with 
high  metallic  lustre ;  it  melts  at  1489°  (Burgess).  Its  density  is  6.8. 
It  is  stable  in  the  air  at  room  temperature,  but  when  strongly  heated 
burns  to  the  oxide  Cr2O3.  It  dissolves  on  warming  with  dilute 
hydrochloric  or  sulphuric  acid,  hydrogen  being  given  off.  Nitric  acid 
does  not  dissolve  it,  but  renders  it  "  passive,"  and  in  this  state  it  i$ 

533 


534    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

not  attacked  by  hydrochloric  or  sulphuric  acid.  Other  oxidizing 
agents  have  the  same  effect  as  nitric  acid  in  rendering  chromium 
inactive  ;  the  nature  of  this  "  passivity  "  is  not  understood  (cf.  iron, 
p.  558).  It  is  remarkable  that  when  dilute  acids  act  on  some  specimens 
of  chromium  the  evolution  of  hydrogen  is  periodic  ;  in  other  words,  a 
period  of  rapid  solution  of  the  metal  alternates  regularly  with  a  com- 
plete cessation  of  action. 

Chromium  is  now  largely  used  in  the  preparation  of  chrome  steel, 
which  contains  up  to  3  per  cent,  of  chromium. 

Compounds  of  Chromium— Chromium  forms  three  series  of 
compounds,  which  may  be  looked  upon  as  being  derived  from  the 
oxides,  CrO,  chromous  oxide,  Cr2O3,  chromic  oxide,  and  CrO3, 
chromic  anhydride. 

Oxide.       Character.  Corresponding  Compounds. 

(CrO)  basic  chromous  salts  ;  example,  CrCl2 

Cr2O3  basic  chromic  salts  „          CrCl3 

Cr03  acidic        {  chromates  ;    „          K2CrO4 

(  dichromates  „          K2Cr2O7 

Chromous  Compounds— The  compounds  of  divalent  chromium  are  powerful 
reducing  agents,  having  a  great  tendency  to  change  into  chromic  salts  by  oxida- 
tion ;  for  this  reason  they  are  very  difficult  to  prepare  in  pure  condition. 

Chromous  chloride,  CrCl2,  is  obtained  in  solution  by  dissolving  chromium 
in  hydrochloric  acid  free  from  oxygen,  and  in  the  anhydrous  condition  as  a  white 
crystalline  compound  by  heating  chromic  chloride  in  a  current  of  hydrogen.  The 
aqueous  solution  of  chromous  chloride  is  blue,  but  rapidly  turns  green  in  the  air 
owing  to  oxidation  : 

4CrCl2+ 4HC1  +  O2->4CrCl3  +  2H2O. 

Chromous  sulphate,  CrSO4,7H2O,  occurs  in  blue  crystals  isomorphous  with 
ferrous  sulphate.  Chromous  acetate,  Cr(C2H3O2)2,  a  red  crystalline  powder 
obtained  by  adding  a  solution  of  chromous  chloride  to  a  saturated  solution 
ot  sodium  acetate,  is  the  most  stable  chromous'  salt.  Chromous  hydroxide, 
Cr(OH)2,  obtained,  by  double  decomposition,  is  yellow  and  rapidly  oxidizes  in  the 
air.  The  corresponding  oxide,  CrO,  has  not  hitherto  been  obtained. 

Chromic  Compounds — Chromic  salts  are  derived  from  chromic 
hydroxide,  Cr(OH)3,  acting  as  a  base,  and  in  practice  can  be  obtained 
by  dissolving  the  base  in  the  appropriate  acid. 

Chromic  Hydroxide,  Cr(OH)3,^H2O,  is  obtained  as  a  bluish- 
green  colloidal  precipitate  (hydrogel)  on  adding  ammonium  hydroxide 
to  a  solution  of  a  chromic  salt.  When  freshly  precipitated  it  dissolves 
in  a  solution  of  sodium  or  potassium  hydroxide  to  form  a  chromite 
(see  below) ;  but  on  boiling  the  solution  it  is  reprecipitated  in  a  less 


CHROMIUM  535 

highly  hydrated,  more  insoluble  form.  When  the  hydroxide  is  heated 
in  air  it  yields  chromic  oxide,  Cr2O3,  a  green  insoluble  powder.  The 
latter  is  used  as  a  pigment  under  the  name  of  chrome-green. 

Chromic  Chloride,  CrCl3,  is  obtained  in  the  anhydrous  form 
as  violet  leaflets  by  heating  a  mixture  of  chromic  oxide  and  charcoal 
in  a  stream  of  chlorine  (cf.  A1C13,  p.  484).  The  anhydrous  chloride  is 
practically  insoluble  in  water  (a  little  goes  into  solution  on  prolonged 
boiling),  but  when  a  trace  of  chromous  chloride  is  added  the  trichlo- 
ride rapidly  dissolves  to  form  a  green  solution.  This  remarkable 
phenomenon  is  not  yet  understood. 

Chromic  chloride  can  also  be  obtained  in  solution  by  dissolving 
chromic  hydroxide  in  hydrochloric  acid,  or  by  reducing  chromates  or 
dichromates  in  hydrochloric  acid  solution.  The  aqueous  solution  is 
usually  green,  and  on  evaporation  green,  deliquescent  crystals  of  the 
formula  CrCl3,6H2O  are  obtained.  When  heated  in  air  the  hydrated 
chloride  is  decomposed,  and  green  chromic  oxide,  Cr2O3,  is  formed  ; 
when,  however,  the  salt  is  heated  in  a  current  of  dry  hydrogen  chlo- 
ride it  yields  the  violet  anhydrous  chloride.  Another  chloride  of 
chromium,  CrCl3,6H2O,  is  obtained  in  blue  crystals  by  dissolving 
chrome  alum  in  hydrochloric  acid  and  passing  hydrogen  chloride 
into  the  solution.  It  dissolves  in  water  to  form  a  bluish-violet  solu- 
tion, and  it  has  been  shown  that  the  chromium  is  present  as  Cr*'ions 
and  all  the  chlorine  is  ionised.  In  the  green  hydrate  only  one  of  the 
chlorine  atoms  is  present  in  the  ionic  form  ;  its  constitution,  according 
to  Werner,  is  represented  by  the  formula  [Cr(H2O)4Cl2]Cl,2H2O  (cf. 
p.  578). 

Chromic  Sulphate,  Cr2(SO4)3,  is  obtained  by  dissolving  freshly 
precipitated  chromic  hydroxide  in  sulphuric  acid  in  the  cold  ;  on 
concentrating  the  violet  solution  in  a  vacuum  violet  crystals,  of  the 
formula  Cr2(SO4)3, 15 H2O,  separate  out.  When  the  violet  solution  is 
heated  it  turns  green  owing  to  hydrolysis,  but  on  standing  for  a  long 
period  at  room  temperature  it  returns  to  the  original  violet  colour. 
As  in  the  case  of  chromic  chloride,  only  the  violet  solution  contains 
the  normal  salt.  The  nature  of  the  green  compound  has  not  been 
conclusively  established  ;  it  may  possibly  be  represented  by  the  formula 
[Cr4O(SO4)4]SO4,  as  only  aboutone-third  of  the  SO4  in  the  green  solution 
is  ionised  (p.  578):  2Cr2(SO4)3-|-H2O->[Cr4O(SO4)4]"SO4//+H2SO4. 

Chrome  Alums — When  an  alkali  sulphate  or  ammonium  sulphate 
is  added  in  the  requisite  proportion  to  a  solution  of  chromic  sulphate 
and  the  solution  is  evaporated  at  a  temperature  not  exceeding  30°, 
a  double  salt  of  the  type  Cr2(SO4)3,MI2SO4,24H2O,  belonging  to  the 


536     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

class  of  alums,  separates  out  in  octahedral  crystals,  which  appeal 
dark  violet  by  reflected,  red  by  transmitted  light.  The  potassium 
salt,  Cr2(SO4)3,K2SO4,24H2O,  which  is  best  known,  is  conveniently 
prepared  by  passing  sulphur  dioxide  into  a  solution  of  potassium 
dichromate,  acidified  with  sulphuric  acid,  the  mixture  being  kept 
cool  throughout  : 

K2Cr2O7  -1-  H2SO4  +  3SO2->Cr2(SO4)3,K2SO4  +  H2O. 

The  aqueous  solutions  of  the  chrome  alums,  like  that  of  chromic 
sulphate,  turn  green  on  heating,  and  slowly  regain  their  original 
colour  at  room  temperature. 

Chromites  —  The  slightly  acidic  character  of  chromic  hydroxide 
is  shown  by  the  fact,  already  mentioned,  that  when  freshly  pre- 
cipitated it  is  soluble  in  potassium  or  sodium  hydroxide  solution. 
The  solution  presumably  contains  an  alkali  chromite,  for  example, 
Cr(OH)2OK  or  CrOOK.  The  naturally  occurring  chrome  iron- 
stone, FeO,Cr2O3,  may  be  regarded  as  ferrous  chromite,  (OO*O)2Fe. 

Chromium  Trioxide,  CrO3,  and  Chromates  —  The  ultimate 
source  of  all  the  chromium  compounds  is  chrome  ironstone, 
FeO,Cr2O3.  The  finely  powdered  ore  is  mixed  with  potassium 
carbonate  and  lime  and  roasted  in  a  reverberatory  furnace.  The 
chief  use  of  the  lime  is  to  keep  the  mass  porous,  and  thus  facilitate 
the  oxidizing  action  of  the  atmospheric  oxygen.  The  products  of  the 
reaction  are  potassium  and  calcium  chromates,  ferric  oxide,  and 
carbon  dioxide  : 


=  6K2CrO4  +  2CaCrO4  +  2  Fe2O3  +  6CO2. 

The  product  is  broken  up  and  extracted  with  water,  whereby 
the  chromates  of  potassium  and  calcium  are  dissolved.  Potassium 
sulphate  is  then  added,  and  by  double  decomposition  potassium 
chromate  and  calcium  sulphate,  the  latter  of  which  is  almost  insoluble 
in  water,  are  formed.  The  solution  is  poured  off  and  treated  with 
sulphuric  acid,  whereby  potassium  dichromate,  K2Cr2O7,  is  formed  : 

2K2CrO4  +  H2SO4->K2Cr2O7  +  K2SO4  +  H2O. 

The  solution  is  then  evaporated,  and  on  cooling  the  dichromate 
separates  in  large  red  triclinic  crystals.  The  object  of  preparing 
the  dichromate  instead  of  the  chromate  is  that  the  former,  being 
much  less  soluble  in  water,  is  more  easily  separated  from  solution  and 
purified  by  recrystallization. 


CHROMIUM  537 

Potassium  Bichromate,  K2Cr2O7,  prepared  as  just  described, 
Is  soluble  about  i  in  10  of  water  at  15°,  and  the  solution  has  a  strongly 
acid  reaction.  When  the  dry  salt  is  heated  it  decomposes,  giving  off 
oxygen  : 

2K2Cr2O7->2K2CrO4  +  Cr2O3+  30. 

When  heated    with    concentrated    sulphuric    acid,    chromium    and 
potassium  sulphates  are  formed  and  oxygen  is  given  off: 


The  same  reaction  takes  place  in  the  presence  of  a  reducing  agent, 
and  therefore  an  acid  solution  of  potassium  dichromate  is  a  powerful 
oxidizing  agent. 

The  equations  expressing  the  oxidizing  action  of  potassium 
dichromate  can  readily  be  written  when  the  salt  is  represented  as 
K2O,2CrO3,  and  it  is  noted  that  2CrO3  give  Cr2O3  and  30  available 
for  oxidation  : 

K2O,2CrO3->K2O  +  Cr2O3  +  30. 

The  equations  representing  the  oxidation  of  sulphur  dioxide  to 
sulphuric  acid  by  potassium  dichromate  may  therefore  be  written  as 
follows  :  — 

K2Cr2O7->K2O  +  Cr2O3+  3O 
K2O  +  Cr2O3  +  4H2SO4->K2SO 
3[S02  +  H20  +  0]->3[H2S04]. 
Adding 


On  the  same  principle  other  equations  representing  the  oxidizing 
action  of  potassium  dichromate  may  readily  be  written.  When 
it  is  heated  with  concentrated  hydrochloric  acid  the  latter  is  oxidized 
to  chlorine  and  water  : 


O  +  3C12. 
Adding 


It  also  acts  as  an  oxidizing  agent  in  the  absence  of  acid.  When 
a  film  of  gelatine  containing  potassium  dichromate  is  exposed  to 
light,  reduction  to  chromic  oxide  takes  place,  and  the  latter  forms 
with  the  gelatine  a  compound  which,  unlike  gelatine  itself,  does  not 


538    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


ItSsolve  or  swell  up  in  warm  water.  This  property  is  taken  advantage 
of  for  photographic  purposes. 

Potassium  dichromate  is  used  as  an  oxidizing  agent  in  batteries 
and  in  the  preparation  of  organic  dyes,  etc.  Sodium  dichromate, 
Na2Cr2O7,2H2O,  which  is  cheaper  and  much  more  soluble  in  water 
than  the  potassium  salt,  is  often  used  in  place  of  the  latter. 
Ammonium  dichromate,  (NH4)2Cr2O7,  has  already  been  mentioned 
(p.  202)  ;  it  gives  free  nitrogen  on  heating  : 

(NH4)2Cr2O7->Cr2O3  +  4H20  +  N2.    k 

When  chromates  or  dichromates  are  treated  with  more  free  acid, 
still  higher  chromates  are  obtained.  Thus  potassium  trichromate, 
K2O,3CrO3  or  K2Cr3O10,  and  potassium  tetrachromate,  K2O,4CrO3  or 
K2Cr4O13,  have  been  isolated. 

Potassium  Chromate,  K2CrO4,  is  obtained  by  adding  the 
requisite  quantity  of  potassium  hydroxide  to  a  solution  of  potassium 
dichromate:  K2Cr2O7  +  2KOH->2K2CrO4  +  H2O.  It  occurs  in  yellow, 
anhydrous  rhombic  crystals,  isomorphous  with  potassium  sulphate. 
It  is  very  soluble  in  water;  the  solution  has  an  alkaline  reaction 
owing  to  hydrolysis. 

A  number  of  chromates  which  are  insoluble  in  water,  and  can 
therefore  be  prepared  by  double  decomposition,  are  of  interest.  Lead 
chromate,  PbCrO4,  a  bright  yellow  powder,  is  used  as  a  pigment 
under  the  name  of  chrome  yellow.  On  treatment  with  sodium  or 
calcium  hydroxide  a  bright  red  basic  salt,  chrome  red,  Pb2O,CrO4, 
is  obtained. 

Barium  Chromate,  BaCrO4,  is  also  used  as  a  pigment. 
Calcium  chromate,  CaCrO4,2H2O,  is  isomorphous  with  gypsum, 
CaSO4,2H2O.  Silver  chromate,  Ag2CrO4,  and  mercurous  chromate, 
Hg2CrO4,  are  red. 

Chromates  are  sometimes  prepared  by  double  decomposition 
between  a  soluble  salt  and  potassium  dichromate.  This  is  owing  to 
the  fact  that  in  the  solution  of  the  latter  salt  CrO4"  ions  are  also 
present  to  a  small  extent  : 


and  when  the  chromate  is  the  less  soluble  salt  it  is  precipitated. 

Chromium  Trioxide,  Chromic  Anhydride,  CrO3—  When 
concentrated  sulphuric  acid  is  added  to  a  cold  concentrated  solution 
of  potassium  dichromate,  after  some  time  chromium  trioxide,  CrO3, 
separates  in  the  form  of  red,  needle-shaped  crystals.  The  crystals 


CHROMIUM  539 

are  very  deliquescent  and  extremely  soluble  in  water,  forming  a  red 
solution  which  probably  contains  dichromic  acid,  H2Cr2O7.  On 
evaporating  the  solution,  however,  the  trioxide  crystallizes  out. 

When  heated  in  air  chromium  trioxide  loses  oxygen,  and  gives 
chromic  oxide,  Cr2O3.  It  is  an  extremely  powerful  oxidizing  agent. 
Warm  alcohol  catches  fire  when  dropped  on  it  ;  it  chars  organic 
matter,  paper,  etc.,  at  once. 

When  a  solution  of  chromic  anhydride  is  neutralized  with  alkali, 
chromates  and  dichromates  are  formed.  The  former,  as  already 
mentioned,  are  of  the  same  type  as  the  sulphates,  the  latter  correspond 
with  the  pyrosulphates,  e.g.  K2Cr2O7,  K2S2O7. 

When  to  an  acidified  solution  of  potassium  dichromate  some 
hydrogen  peroxide  is  added  and  the  mixture  is  shaken  up  with  ether, 
it  will  be  observed  that  the  layer  of  ether  which  separates  out  is 
coloured  deep  blue.  If  no  ether  is  used,  the  blue  colour  observed 
when  the  hydrogen  peroxide  is  added  rapidly  disappears,  oxygen 
being  given  off,  and  the  substance  giving  rise  to  it  is  therefore  very 
unstable  (p.  142).  The  blue  compound  has  not  been  definitely  isolated, 
but  is  believed  to  be  a  perchromic  acid,  of  the  formula  HCrO6. 

Chromyl  Chloride,  CrO2Cl2,  is  prepared  by  heating  a  mixture 
of  potassium  dichromate,  sodium  chloride  and  concentrated  sul- 
phuric acid  ;  the  chromyl  chloride  distils  over  as  a  deep  red,  fuming 
liquid  which  boils  at  118°.  It  is  readily  decomposed  by  water  into 
chromic  acid  and  hydrochloric  acid  : 


Chromyl  chloride  corresponds  with  sulphuryl  chloride,  SO2C12,  and 
may  be  regarded  as  being  derived  from  chromic  acid,  CrO2(OH)2,  by 
substituting  two  Cl  atoms  for  the  two  OH  groups.  The  intermediate 
compound,  CrO2(OH)Cl,  chlorochromic  acid,  is  not  known,  but  its 
potassium  salt,  potassium  chlorochromate,  CrO2(OK)Cl,  is  obtained 
in  yellowish-red  crystals  by  crystallizing  potassium  dichromate  from 
a  concentrated  solution  in  hydrochloric  acid  : 

K2Cr2O7  +  2HCl<pCrO2(OK)Cl  +  H2O. 

Chromium  Ammonia  Compounds  —  From  chromic  chloride, 
by  the  action  of  ammonia,  a  number  of  complex  compounds  have 
been  derived,  of  the  empirical  formulae  Cr(NH3)6Cl3,  Cr(NH3)5Cl3, 
Cr(NH3)4Cl3,  etc.,  which  behave  like  the  better  known  cobaltammines, 
of  analogous  constitution.  It  is  therefore  more  convenient  to  deal 
with  these  compounds  at  a  later  stage  (pp.  567,  577). 


540    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Tests  for  Chromium — Compounds  of  chromium  are  readily 
recognized  by  their  characteristic  colours.  Chromic  salts  give  with 
alkali  hydroxides  a  green  precipitate  of  chromium  hydroxide,  soluble 
m  excess  of  alkali,  but  reprecipitated  on  boiling.  When  fused  with 
alkali  and  an  oxidizing  agent,  chromic  salts  yield  chromates  (p.  536) 
which  are  characterized  by  a  yellow  colour,  and  formation  of  char- 
acteristic precipitates  with  soluble  lead  and  silver  salts.  Further, 
both  chromates  and  dichromates  act  as  oxidizing  agents,  especially 
in  acid  solution,  being  thereby  reduced  to  green  chromic  salts. 

MOLYBDENUM 

This  comparatively  rare  element  occurs  chiefly  as  molybdenite,  MoS2,  which 
has  a  strong  resemblance  to  graphite,  and  as  wulfenite,  PbMoO4.  When  molyb- 
denite is  roasted,  it  is  converted  to  the  trioxide,  MoO3,  which  is  a  white  powder. 
Molybdenum  can  be  prepared  from  the  trioxide  by  heating  in  a  current  of  hydro- 
gen, or  by  heating  with  carbon  in  the  electric  furnace. 

Molybdenum  is  a  hard  metal,  of  density  9.1,  and  strongly  resembles  iron; 
like  the  latter  metal  it  takes  up  carbon  and  may  be  welded  and  tempered. 

The  most  important  oxide  of  molybdenum  is  the  trioxide,  MoO3,  which,  like 
chromium  trioxide,  CrO3,  is  exclusively  acidic.  It  readily  dissolves  in  solutions  of 
alkali  hydroxide  or  ammonium  hydroxide  to  form  molybdates — examples,  sodium 
molybdate,  Na2MoO4,ioH2O,  and  ammonium  molybdate,  (NH4)2MoO4.  When 
a  strong  acid  is  added  to  a  molybdate  solution,  molybdic  acid,  H2MoO4,H2O, 
separates  in  colourless,  lustrous  crystals  ;  it  is  soluble  in  excess  of  acid.  Ammonium 
molybdate,  dissolved  in  excess  of  moderately  concentrated  nitric  acid,  is  used  as  a 
test  for  phosphates.  When  the  phosphate  is  added  to  the  molybdate  reagent  and  the 
mixture  warmed,  ammonium  phosphomolybdate,  of  the  approximate  composition 
(NH4);>PO4,iiMoO3,6H2O,  comes  down  as  a  yellow  precipitate.  The  composition 
of  this  precipitate  indicates  the  tendency  to  the  formation  of  polymolybdates, 
which  is  much  more  pronounced  'than  the  tendency  to  form  polychromates 
(P-  538). 

Besides  the  trioxide,  the  compounds  Mo2O3  and  MoO2  are  definitely  known ; 
neither  of  them  is  definitely  basic.  Chlorides  of  the  formulae  (MoCl2)3 ;  MoCl3  ; 
MoCl4  and  MoCl3  have  been  described. 

TUNGSTEN  OR  WOLFRAM 

Tungsten  is  found  in  the  relatively  scarce  minerals,  wolfram,  Fe(Mn)WO4,  and 
scheelite,  CaWO4.  The  element  itself  is  made  by  reducing  tungstic  acid,  H2WO4, 
with  aluminium  according  to  the  Goldschmidt  process.  Tungsten  is  a  very  hard 
steel-gray  metal  of  density  about  19 ;  it  is  stable  in  the  air  under  ordinary  con- 
ditions, but  on  heating  forms  the  trioxide.  When  added  to  steel  in  the  propor- 
tion of  about  5  per  cent.,  a  very  hard  alloy  (tungsten  steel)  is  formed.  Tungsten 
is  also  used  in  making  filaments  for  incandescent  lamps. 

On  fusing  tungsten  with  sodium  carbonate  and  extracting  with  water,  a  solution 
of  sodium  tungstate,  Na2WO4,  is  obtained.  By  adding  acid  to  the  solution, 
tungstic  acid,  H2WO4,H2O,  is  obtained;  when  the  latter  is  ignited  tungsten 


ELEMENTS   OF  THE   CHROMIUM   SUB-GROUP    541 

trioxide,  WO3,  remains.  Like  molybdenum  trioxide,  tungsten  trioxide  has  a 
considerable  tendency  to  unite  with  other  compounds,  forming  complexes  con- 
taining a  large  number  of  WO3  groups.  Phosphotungstic  acid,  which  is  used  as 
a  reagent,  has  the  formula  H3PO4,*WO3,  where  x  is  a  large  number,  probably 
variable  according  to  the  conditions  of  preparation. 

Four  tungsten  chlorides,  WC12,  WC14,  WC15,  WC16,  and  two  oxychlorides, 
WO2C12  and  WOC14,  are  known.  Tungsten  hexachoride,  WC16>  prepared  by 
the  direct  action  of  chlorine  on  tungsten,  occurs  in  dark  violet  crystals,  which 
melt  at  275°. 

URANIUM 

The  chief  source  of  uranium  compounds  is  pitch-blende  or  uraninite,  which  is 
found  in  Austria  and  in  Cornwall ;  it  consists  mainly  of  the  oxide,  UO2'2UO3 
or  U3O8.  In  recent  years  pitch-blende  has  become  very  important  owing  to  the 
fact  that  it  contains  radium  (p.  580).  Uranium  is  also  of  interest  as  being  the 
element  with  the  highest  atomic  weight  (238.5). 

The  metal  is  obtained  by  heating  the  chloride  with  sodium  or  by  reducing  the 
oxide  with  carbon  in  the  electric  furnace.  It  is  a  hard  silvery-white  metal ;  its 
density  is  18.7. 

There  are  five  oxides  of  uranium  :  UO2,  which  is  exclusively  basic  ;  U2O3  ;  U3O8, 
which  is  the  most  stable  in  the  air  ;  UO8,  which  is  both  basic  and  acidic,  and 
UO4,  which  is  a  peroxide.  There  are  two  chief  series  of  salts.  The  uranous 
salts,  in  which  the  uranium  is  quadrivalent,  are  obtained  by  dissolving  the  dioxide 
in  strong  acids  ;  salts  of  this  type  are  nranous  chloride,  UC^,  and  uranous 
sulphate,  U(SO4)2,8H2O. 

The  second  series  of  salts,  which  contain  hexavalent  uranium,  are  derived  from 
the  compound  UO2(OH)2,  which  is  usually  called  uranic  acid.  They  thus  con- 
tain the  divalent  group  UO2,  the  so-called  uranyl  group.  Among  the  salts 
belonging  to  this  class  are  uranyl  nitrate,  UO2(NO3)2,6H2O ;  uranyl  sulphate, 
UO2SO4,3H2O ;  and  uranyl  chloride,  UO2C12.  The  salts  are  prepared  by  dis- 
solving uranic  acid  in  the  appropriate  acid. 

As  its  name  implies,  the  compound  UO2(OH)2  has  also  acidic  properties.  The 
salts  obtained  by  treating  it  with  alkali  hydroxides  are,  however,  diuranates 
analogous  in  constitution  to  dichromates.  Sodium  diuranate,  Na2U2O7,6H2O, 
which  is  called  uranium  yellow,  is  used  for  tinting  glass. 

Two  uranium  chlorides,  UC14  and  UC15,  are  definitely  known. 

Summary  of  Chromium  Sub-group — As  chromium  is  the 
only  important  member  of  the  sub-group,  it  is  unnecessary  to  com- 
pare the  properties  of  the  elements  in  detail.  They  are  characterized 
by  forming  several  classes  of  compounds  with  different  valencies, 
but  in  the  most  important  compounds  they  are  hexavalent,  corre- 
sponding with  their  position  in  the  periodic  table.  The  oxides  CrO3, 
MoO3,  WO3  and  UO3  are  almost  exclusively  acidic  ;  the  best  known 
compounds  derived  from  them  are  of  the  types  K2Cr(Mo,W,Ur)O4  and 
K2Cr2(Mo2,W2,Ur2)O7.  There  is  a  considerable  tendency  throughout 
to  form  complex  salts  containing  a  number  of  EO3  groups.  Sulphur 


542     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

also  shows  this  tendency  to  a  slight  extent  (cf.  pyrosulphuric  acid, 
H2S207). 

As  regards  analogies  with  members  of  sub-group  A,  the  four  elements 
now  under  consideration  are  most  closely  allied  with  sulphur.  Many 
of  their  compounds  are  of  the  same  type  as  those  of  sulphur,  and  in 
some  cases  are  isomorphous  with  them.  This  is  shown  in  K2CrO4 
and  K2SO4,  and  in  Na2CrO4,ioH2O  and  Na2SO4,ioH2O.  The 
analogy  of  sulphuryl  chloride,  SO2C12,  with  chromyl  chloride,  CrO2Cl2, 
has  already  been  referred  to. 

Chromium  also 'shows  analogies  with  metals  belonging  to  other 
groups.  The  chromic  compounds,  for  example,  show  similarity  with 
those  of  aluminium  and  of  trivalent  iron,  as  is  evident  from  the 
existence  of  the  isomorphous  alums.  Further,  there  is  some  analogy 
between  compounds  of  divalent  chromium  and  divalent  iron,  e.g. 
CrSO4,7H2O  and  FeSO4,7H2O  are  isomorphous.  Further,  an  oxide 
of  iron  of  the  formula  FeO3  is  known  (p.  560). 


CHAPTER  XXXIV 

THE  MANGANESE  SUB-GROUP  (GROUP  VII, 
SUB-GROUP  A) 

MANGANESE  and  the  halogens  are  the  only  known  members 
of  the  seventh  group.     The  former  element,  like  those  belong- 
ing to  the  sixth  group  just  considered,  forms  compounds  belonging  to 
a  variety  of  stages  of  oxidation,  and  shows  analogies  with  magnesium, 
iron,  chromium  and  other  elements,  as  well  as  with  the  halogens. 

MANGANESE 

Symbol,  Mn.     Atomic  Weight— 54.93. 

Occurrence — Manganese  does  not  occur  naturally  in  the  free 
condition  (except  occasionally  in  meteorites),  but  in  combination  with 
oxygen  it  occurs  fairly  abundantly  as  pyrolusite,  MnO2 ;  also  as 
hausmannite,  Mn3O4,  braunite,  Mn2O3,  manganite,  Mn2O3,H2O, 
and  manganese  spar,  MnCO3. 

Preparation — The  metal  is  readily  obtained  pure  by  reducing 
the  dioxide,  MnO2,  with  aluminium  according  to  the  Goldschmidt 
process. 

Properties — Manganese  is  a  reddish-gray  lustrous  brittle  metal, 
harder  than  iron.  Its  density  is  8.0,  and  it  melts  at  1207°.  It  becomes 
superficially  oxidized  in  moist  air  and  decomposes  water  at  100°,  with 
evolution  of  hydrogen.  It  dissolves  readily  in  acids,  even  in  weak 
acids  such  as  acetic  acid,  hydrogen  being  given  off  and  manganous 
salts  formed. 

Manganese  is  largely  used  as  an  addition  to  iron  (p.  558),  and  also 
in  the  production  of  manganese  bronze  (p.  427). 


OXIDES  OF  MANGANESE 


OXIDES  OF  MANGANESE 

Six  oxides  of  manganese  are  known :  manganous  oxide,  MnO  ;  di- 
manganese  trioxide,  Mn2O3  ;  manganese  dioxide,  MnO2;  trimanganest 
tetroxide,  Mn3O4;  manganese  trioxide,  MnO3,  and  manganese  hept- 
oxide,  Mn2O7. 


544     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Oxide.  Character.                        Corresponding  Compounds. 

MnO  basic  Manganous  Salts  ;  example  MnCl2 

Mn3O4  basic  (mixed  oxide) 

Mn2O3  basic  Manganic  Salts                „      MnCl3 

MnO2  acidic  and  peroxide  Manganites                       „      CaMnO3 

MnO3  acidic  Manganates                      „      K2MnO4 

Mn2O7  acidic  Permanganates                 „      KMnO4 

Manganous  Oxide,  MnO,  is  prepared  by  heating  manganotis 
carbonate,  MnCO3,  in  absence  of  air,  or  by  heating  any  of  the  other 
oxides  in  a  current  of  hydrogen.  It  is  a  green  powder,  which  oxidizes 
to  Mn3O4  on  heating  strongly  in  air. 

Manganous  Hydroxide,  Mn(OH)2,  is  obtained  as  a  white 
precipitate  on  mixing  air-free  solutions  of  an  alkali  hydroxide  and  a 
manganous  salt.  It  rapidly  absorbs  oxygen  from  the  air,  forming 
the  green  hydroxide,  Mn(OH)3. 

Trimanganese  Tetroxide,  Mn3O4,  occurs  naturally  in  red 
prismatic  crystals  as  hausmannite,  and  is  obtained  on  heating  any 
of  the  other  oxides  of  manganese  in  the  air  at  1000°.  When  it  is  heated 
with  dilute  sulphuric  acid,  manganous  sulphate  goes  into  solution  and 
the  dioxide  remains  (cf.  Pb3O4,  p.  503)  : 


When  it  is  treated  with  cold  concentrated  sulphuric  acid,  a  mixture 
of  manganous  and  manganic  sulphates  is  probably  formed. 

Manganic  Oxide,  Mn2O3,  occurs  naturally  as  braunile,  and 
is  obtained  by  heating  any  of  the  other  oxides  in  oxygen  at  650- 
900°.  When  heated  with  dilute  sulphuric  acid,  manganous  sulphate 
dissolves  and  manganese  dioxide  remains  insoluble.  For  this  reason 
it  is  sometimes  represented  as  MnO'MnO2  ;  on  the  other  hand, 
corresponding  salts,  the  manganic  salts,  are  known. 

Manganese  Dioxide,  MnO2,  occurs  naturally  in  grayish- 
black  crystals  as  pyrolusite.  It  is  best  obtained  in  a  pure  con- 
dition by  cautiously  heating  manganous  nitrate.  When  it  is 
heated  with  hydrochloric  acid,  chlorine  is  given  off  and  manganous 
chlorine,  MnCl2,  remains  in  solution  : 


This  method  is  used  for  the  commercial  preparation  of  chlorine,  but 
is  now  to  some  extent  replaced  by  electrolytic  methods  (p.  399). 
When  the  dioxide  is  treated  with  cold  concentrated  hydrochloric 


MANGANESE 


545 


acid,  it  dissolves  to  a  dark  liquid  and  very  little  chlorine  is  given  off; 
on  warming  the  solution  chlorine  is  given  off  and  manganous  chloride 
remains  in  solution.  The  nature  of  the  compound  or  compounds 
present  in  the  cold  solution  is  not  settled.  It  may  contain  the  tetra- 
chloride,  MnCl4,  the  trichloride,  MnCl3,  or  a  mixture  of  the  two. 

When  manganese  dioxide  is  heated  with  concentrated  sulphuric 
acid,  manganous  sulphate  is  formed  and  oxygen  is  liberated  : 

2MnO2+2H2SO4->2MnSO4+2H2O  +  O2. 

Manganese  dioxide  is  used  in  glass-making,  and  as  an  oxidizing 
agent  in  galvanic  batteries. 

Manganous  Acid  (Hydrated  Manganese  Dioxide),  Mn(OH)4, 
is  obtained  as  a  black  precipitate  by  the  action  of  a  hypochlorite  on 
a  manganous  salt  in  a  neutral  or  alkaline  medium.  On  partial 
dehydration,  the  compound  MnO(OH)2  or  H2MnO3,  a  dibasic 
manganous  acid,  is  obtained. 

The  Weldon  process  for  the  utilization  of  the  manganous  chloride 
formed  in  the  preparation  of  chlorine  from  pyrolusite  is  based  on 
the  formation  of  salts,  manganites,  derived  from  the  dibasic  acid. 
Milk  of  lime  is  added  to  the  chloride  solution  and  air  forced 
through  the  mixture,  whereby  the  manganous  hydroxide  first 
precipitated  is  converted  into  calcium  manganite,  CaMnO3  (or 
CaO.MnO2): 

MnCl2  +  2CaO  +  O->CaMnO3+  CaCl2. 

The  manganite,  being  insoluble  in  water,  slowly  settles  as  a  black 
mud,  the  calcium  chloride  solution  is  poured  off,  and  the  manganite 
treated  with  hydrochloric  acid  for  the  production  of  more  chlorine  : 


Manganese  Trioxide,  MnO3,  prepared  by  adding  potassium 
permanganate,  dissolved  in  sulphuric  acid,  to  dry  sodium  carbonate, 
occurs  as  a  dark  red  mass.  It  is  the  anhydride  of  manganic 
acid  (q.v.\ 

Dimanganese  Heptoxide,  Mn2O7,  is  obtained  as  a  green, 
oily  liquid  by  cautiously  adding  potassium  permanganate  to  con- 
centrated sulphuric  acid,  the  mixture  being  well  cooled.  The 
heptoxide  is  very  unstable,  and  decomposes  violently  on  warming, 
forming  the  dioxide  and  oxygen.  It  is  the  anhydride  of  perman- 
ganic acid,  HMnO4  (q.v.\ 
35 


546     A    TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


MANGANOUS  SALTS 

All  the  oxides  of  manganese  on  heating  with  acids  yield  manganous 
salts,  in  which  the  metal  is  divalent.  When  the  higher  oxides  are 
used,  either  oxygen  is  given  off  or  the  acid  is  oxidized.  The 
manganous  compounds  are  the  only  stable  salts  containing  Mn 
ions.. 

Manganous  Chloride,  MnCl2,  obtained  by  dissolving  any  of 
the  oxides  in  hydrochloric  acid,  separates  from  solution  in  pink  crystals 
with  4H2O.  It  can  be  obtained  in  the  anhydrous  form  by  heating 
the  tetrahydrate  in  a  current  of  hydrogen  chloride,  or  by  heating 
manganous  ammonium  chloride,  MnCl2,2NH4Cl,H2O  (cf.  magnesium 
chloride,  p.  463). 

Manganous  Sulphate,  MnSO4,  prepared  by  the  general 
method,  separates  from  solution  below  6°  in  pink,  monoclinic 
crystals  isomorphous  with  ZnSO4,7H2O ;  between  6°  and  20°  it 
is  obtained  as  MnSO4,5H2O  in  triclinic  crystals  isomorphous 
with  CuSO4,5H2O  ;  above  20°,  MnSO4,4H2O  separates  in  rhombic 
prisms.  It  forms  double  salts  with  alkali  sulphates,  of  the  type 
MnSO4,K2SO4,6H2O,  isomorphous  with  the  similarly  constituted 
salts  of  magnesium,  zinc  and  iron  (p.  468). 

Manganous  Sulphide,  MnS,  is  obtained  as  a  green  powder 
by  heating  any  of  the  oxides  in  a  stream  of  hydrogen  sulphide  ;  and 
in  the  hydrated  form  as  a  pink  precipitate  when  ammonium  sulphide 
is  added  to  a  solution  of  a  manganous  salt.  It  is  soluble  in  dilute 
acids,  even  in  acetic  acid. 


MANGANIC  SALTS 

Manganic  Hydroxide,  Mn(OH)3,  from  which  the  manganic 
salts  are  derived,  has  already  been  mentioned.  It  is  an  extremely 
weak  base,  and  the  salts  are  practically  completely  hydrolyzed  on 
addition  of  water. 

Manganic  Chloride,  MnCl3,  is  probably  formed  by  the  action 
of  cold  hydrochloric  acid  on  manganese  dioxide  (p.  544),  but  has  not 
been  isolated  from  the  solution.  It  has,  however,  been  obtained  by 
treating  manganese  dioxide,  suspended  in  carbon  tetrachloride,  with 
dry  hydrogen  chloride,  and  then  extracting  the  trichloride  by  means 
of  ether.  It  is  a  nearly  black  solid  with  a  gree*nish  tinge,  and  is 
immediately  decomposed  by  water. 


MANGANESE  547 

Manganic  Sulphate,  Mn2(SO4)3,  is  obtained  as  a  dark-green 
powder  by  gently  heating  hydrated  manganese  dioxide  with  con- 
centrated sulphuric  acid.  It  is  immediately  hydrolyzed  by  water. 
With  alkali  sulphates  it  forms  alums,  e.g.  Mn2(SO4)3,K2SO4,24H2O, 
which  are  more  stable  than  the  salt  itself. 

MANGANATES  AND  PERMANGANATES 

Manganates  —  The  manganates  are  derived  from  manganic  acid, 
H2MnO4  (or  MnO2(OH)2),  which  has  not  itself  been  isolated,  but  the 
oxide,  MnO3,  is  known  (p.  545).  Potassium  manganate  is  obtained  as 
a  green  mass  by  fusing  manganese  dioxide  with  potassium  hydroxide 
or  carbonate  with  or  without  the  addition  of  an  oxidizing  agent  such 
as  potassium  nitrate  or  chlorate.  When  no  oxidizing  agent  is  added 
the  oxidation  is  effected  by  oxygen  taken  up  from  the  air: 


The  fused  mass  is  extracted  with  water  and  the  solution  evaporated, 
when  potassium  manganate  is  deposited  in  green  rhombic  crystals, 
isomorphous  with  potassium  sulphate  and  potassium  chromate. 

Potassium  manganate  is  stable  only  in  the  presence  of  alkali  ;  when 
the  solution  is  diluted  and  warmed,  still  more  readily  when  a  weak 
acid  is  added  (for  example,  when  carbon  dioxide  is  passed  into  the 
solution),  the  colour  changes  to  deep  purple,  owing  to  the  formation 
of  potassium  permanganate,  manganese  dioxide  being  precipitated  : 


The  remarkable  reaction  just  considered  is  in  the  nature  of  a 
hydrolysis,  and  as  the  oxide  corresponding  with  potassium  perman- 
ganate is  Mn2O?,  and  MnO2  is  also  formed,  it  is  clear  that  a  simultaneous 
oxidation  and  reduction  of  the  manganate  (whose  corresponding  oxide 
is  MnO3)  has  occurred.1  We  may  assume  that  in  the  first  stage  of  the 
hydrolysis,  the  very  unstable  manganic  acid  is  formed  : 

3x[K2MnO4  +  2H2O->H2MnO4  +  2KOH]    (i), 

and  immediately  decomposes  into  permanganic  acid,  HMnO4,  and 
manganese  dioxide  : 

2H2O     (2), 


1  Other  instances  of  the  decomposition  of  a  compound  into  two  others,  one  in 
a  higher  and  one  in  a  Icwer  stage  of  oxidation  than  the  original  substance,  have 
already  been  given  (cf.  potassium  chlorate,  p.  181). 


548     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

the  permanganic  acid   finally  uniting   with   part   of  the  potassium 
hydroxide  to  form  potassium  permanganate  : 


4  +  KOH-»KMn04+H20]. 
Adding 


As  Ostwald  has  pointed  out,  it  is  simpler  to  write  the  reaction 
in  terms  of  ions  : 


This  shows  that  H'  ions  are  used  up  in  the  change,  which  therefore 
can  only  proceed  in  acid  or  nearly  neutral  solution.  In  alkaline 
solution,  practically  no  H'  ions  are  present,  and  therefore  the 
manganate  is  stable.1 

Manganates  act  as  oxidizing  agents  in  alkaline  solution,  being 
thereby  reduced  to  manganese  dioxide  (MnO3->MnO2+O). 

Permanganic  Acid  and  Permanganates—  The  perman- 
ganates are  derived  from  permanganic  acid,  HMnO4;  the  solutions 
contain  the  MnO/  ion,  which  is  purple. 

Permanganic  Acid,  HMnO4,  is  obtained  by  adding  to  a  solu- 
tion of  barium  permanganate  the  requisite  quantity  of  sulphuric  acid, 
and  filtering  from  the  precipitated  barium  sulphate  ;  on  evaporating 
the  solution,  the  acid  is  obtained  in  violet-blue  crystals.  The  acid  is 
also  obtained  in  aqueous  solution  by  dissolving  manganese  heptoxide, 
Mn2O7,  in  cold  water. 

Permanganic  acid  is  a  strong  monobasic  acid,  comparable  in 
strength  with  hydrochloric  acid,  and  is  a  powerful  oxidizing  agent. 

Potassium  Permanganate,  KMnO4,  prepared  as  described 
above,  occurs  in  purple  rhombic  prisms,  isomorphous  with  potassium 
perchlorate.  The  aqueous  solution  of  the  salt  is  deep  purple.  At 
20°  100  grams  of  water  dissolve  6.35  grams,  at  50°  16.9  grams  of  the 
salt. 

Potassium  permanganate  is  a  powerful  oxidizing  agent.  When  the 
dry  salt  is  heated,  it  decomposes  thus  : 

2KMnO4->K2MnO4+  MnO2+  O2. 
When  it  is  heated  in  solution  with  an  alkali,  the  reaction  by  which  it 

1  In  the  conversion  of  the  MnO4"  to  the  MnO4'  ion  there  is  a  diminution  in  the 
number  of  negative  charges,  and  the  reaction  therefore  conforms  to  the  extended 
definition  of  oxidation  already  given  (p.  431). 


MANGANESE  549 

is  formed  is   reversed,  and  a  green  solution  of  the  manganate  is 
obtained  : 


When  a  reducing  agent  is  present  in  alkaline  solution,  further  reduc- 
tion to  the  dioxide  occurs,  so  that  from  two  molecules  of  permanganate 
in  alkaline  solution,  three  atoms  of  oxygen  are  available  for  oxidation 


In  acid  solution,  in  the  presence  of  a  reducing  agent,  reduction 
proceeds  to  manganous  salt,  so  that  from  two  molecules  of  per- 
manganate, five  atoms  of  oxygen  are  available  for  oxidation 
(Mn2O7->2MnO  +  $0)  : 


In  the  light  of  the  explanations  given  in  previous  cases,  there 
should  be  no  difficulty  in  writing  the  equations  representing  reactions 
of  this  type.  Suppose,  for  instance,  it  is  required  to  write  the  equa- 
tion representing  the  oxidation  of  ferrous  sulphate,  FeSO4,  to  ferric 
sulphate,  Fe2(SO4)3,  in  the  presence  of  free  sulphuric  acid.  As  the 
oxide  corresponding  with  ferrous  sulphate  is  FeO,  and  that  corre- 
sponding with  ferric  sulphate  is  Fe2O3,  it  is  clear  that  one  atom  of 
oxygen  will  oxidize  two  molecules  of  ferrous  to  one  of  ferric  salt.  We 
have  therefore 


(i) 
5  x  [2FeS04  +  H2S04  +  0-»Fe2(S04)3+H20].  (2) 

Adding 

ioFeS04  +  8H2S04+2KMn04-»5Fe2(S04)3  +  K2S04  + 
2MnSO4  +  8H2O. 

Sulphur  dioxide,  hydrogen  sulphide,  hydrogen  peroxide,  nitrous 
acid  and  other  compounds  reduce  potassium  permanganate  to  man- 
ganous salts  in  acid  solution,  and  the  equations  can  be  written  as  in 
the  above  example.  The  partial  equation  (i)  is  the  same  in  each 
case.  Further,  as  manganous  salts  are  almost  colourless  in  solution, 
potassium  permanganate  can  be  used  for  the  quantitative  estimation 
of  many  readily  oxidizable  substances.  If  the  permanganate  is  added 
to  the  reducing  agent,  the  end  of  the  reaction  is  marked  by  the 
appearance  of  a  permanent  pink  colour  due  to  free  permanganate, 
and  no  other  indicator  is  required. 


550    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

On  account  of  their  powerful  oxidizing  action,  permangates  are 
largely  used  as  disinfectants.  The  cheaper  sodium  permanganate, 
which  does  not  readily  crystallize,  is  sold  in  solution  for  this  purpose 
under  the  name  of  Condy's  disinfecting  fluid.  The  action  is  identical 
with  that  of  the  potassium  salt,  being  due  to  the  MnO4'  ion. 

Tests  for  Manganese— The  formation  of  a  flesh-coloured 
sulphide,  MnS,  when  ammonium  sulphide  is  added  to  the  solution 
of  a  manganous  salt  is  a  characteristic  test ;  the  sulphide  is  soluble  in 
acetic  acid.  Further,  the  formation  of  a  green  mass  of  manganate 
when  a  salt  of  manganese  is  fused  with  alkali  and  an  oxidizing  agent 
is  a  useful  test ;  when  the  green  mass  is  treated  with  water  and  the 
solution  is  boiled,  it  becomes  pink.  Manganese  salts  give  an  amethyst 
colour  to  a  borax  bead. 

Summary — Owing  to  the  isolated  position  of  manganese  'in 
the  periodic  table,  there  are  no  elements  which  very  closely  resemble 
it  ;  but,  on  the  other  hand,  it  shows  a  number  of  points  of  analogy 
with  elements  belonging  to  other  groups.  The  remarkable  variety  of 
valencies  shown  by  manganese  has  already  been  mentioned ;  it  acts 
as  a  divalent,  trivalent,  hexavalent,  septavalent,  and  possibly  quadri- 
valent element.  In  the  divalent  manganous  compounds  it  re- 
sembles magnesium  and  ferrous  salts,  the  compounds  MnSO4,7H2O, 
MgSO4,yH2O  and  FeSO4,7H2O  being  isomorphous;  in  its  trivalent 
compounds  it  resembles  iron  and  aluminium,  as  is  clearly  shown  by 
the  existence  of  a  manganese  alum,  Mn2(SO4)3,K2SO4,24H2O  ;  in  its 
sexavalent  compounds  it  resembles  sulphur  and  chromium,  as  is 
shown  by  the  isomorphism  of  the  compounds  K2MnO4,  K2SO4  and 
K2CrO4  ;  finally,  in  its  septavalent  compounds  it  resembles  the  highest 
oxidation  stage  of  chlorine,  the  compounds  KMnO4  and  KC1O4  being 
isomorphous.  It  follows  that  the  formula  of  potassium  manganate 

O^          /OK 
is  represented  graphically  thus :    •.  ^Mn<^        ,  and  potassium  per- 


,>0 
manganate  thus : 

O^          \OK. 

It  is  interesting  to  note  that  the  green  manganate  ion  and  the  red 
permanganate  ion  are  both  of  the  composition  MnO4,  but  the  former 
has  two  negative  charges,  the  latter  only  one. 


CHAPTER  XXXV 

THE  IRON  SUB-GROUP  (GROUP  VIII) 
Iron,  55.85.     Cobalt,  58.97.     Nickel,  58.68. 

THE  members  of  the  iron  family  of  elements,  unlike  the  other 
families  of  elements  already  considered,  belong  to  the  same 
period,  and  their  atomic  weights  are  not  very  different.  As  already 
explained,  they  are  the  middle  members — sometimes  known  as  the 
transitional  members— of  the  first  long  period,  and  belong  to  the 
so-called  eighth  group. 

Iron  forms  two  series  of  compounds,  ferrous  compounds,  in  which 
it  is  bivalent,  &&&  ferric  compounds,  in  which  it  is  trivalent.  Cobalt 
also  functions  as  a  bivalent  and  trivalent  element,  but  nickel  forms 
only  one  series  of  salts,  in  which  it  is  bivalent. 

IRON 

Symbol,  Fe.     Atomic  weight= 55.85. 

History — Iron  has  been  known  from  very  early  times ;  it  is 
mentioned  both  in  Homer  and  in  the  Pentateuch.  In  the  earliest 
stage,  iron  weapons  were  probably  made  of  meteoric  iron,  but  the 
ancient  Egyptians  and  Assyrians  were  familiar  with  the  preparation 
of  iron  from  its  ores.  The  transition  from  the  bronze  age  to  the  iron 
age  marked  a  great  advance  in  human  culture. 

Occurrence — Iron  is  a  very  abundant  and  a  very  widely  distri- 
buted element.  In  the  free  condition,  it  is  met  with  in  meteorites, 
which  also  contain  cobalt,  nickel,  chromium,  manganese  and  other 
elements  in  smaller  proportion. 

The  chief  ores  of  iron  used  in  preparing  the  metal  are :  red  hcematite 
and  specular  iron  ore,  Fe2O3  ;  brown  hcematite,  2Fe2O3,3H2O  ;  mag- 
netic iron  ore,  Fe3O4  ;  spathic  iron  ore,  FeCO3  ;  clay  ironstone,  which 
is  ferrous  carbonate  mixed  with  clay,  and  black  band,  which  consists 
of  ferrous  carbonate  mixed  with  clay  and  coal.  It  also  occurs  in  iron 
pyrites,  FeS2,  but  these  are  used  only  as  a  source  of  sulphur. 

Iron  is  met  with  in  most  rocks  and  in  all  soils,  and  is  an  essential 
constituent  for  the  growth  of  plants. 

5S« 


552     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Metallurgy  of  Iron— On  the  large  scale,  iron  is  reduced  from 
its  ores  by  means  of  carbon  at  a  high  temperature.  The  arrange- 
ment used  for  this  purpose,  known  as  a  blast-furnace,  is  illustrated  in 
Fig.  89.  Such  a  furnace  is  from  50  to  100  feet  high  and  14  to  20  feet 

wide  at  the  broadest  part  ;  it 
is  built  of  masonry  and  lined 
with  firebrick.  The  charge, 
which  consists  of  the  ore  along 
with  coke  or  coal  and  lime- 
stone, is  admitted  at  the  top 
by  the  so-called  cup  and  cone 
arrangement.  The  cone,  «, 
which  normally  presses  against 
the  funnel  and  thus  prevents 
the  furnace  gases  from  escap- 
ing, is  momentarily  lowered  to 
admit  the  charge.  At  the 
base  of  the  furnace,  hot  air  is 
forced  in  by  the  pipes  b,b> 
known  as  tuyeres.  The  fur- 
nace gases,  which  are  rich  in 
carbon  monoxide,  are  led  oft 
by  the  pipe,  c,  and  are  used  for 
heating  the  air  for  the  blast, 
and  for  driving  gas-engines. 

The  chemical  changes  taking 
place  in  the  furnace  will  now 
be  discussed.  In  the  upper 
part  of  the  furnace,  water  and 
carbon  dioxide  are  expelled l 
and  ferrous  carbonate  is  con- 
verted into  ferric  oxide  ;  the 
limestone  is  simultaneously 
changed  to  lime  (calcium 
oxide).  As  the  air  forced  in  at 
the  base  of  the  furnace  is  at  a 


FIG.  89. 


very  high  temperature  (about  800°)  it  combines  with  the  carbon, 
forming  carbon  dioxide.  The  latter,  on  its  way  upwards  through  the 
hottest  part  of  the  furnace  at  /  is  reduced  by  the  red-hot  carbon  to 

1  Some  ores  are  calcined  before  being  introduced  into  the  furnace,  the  object 
being  to  expel  water  and  carbon  dioxide  and  render  them  more  porous. 


IRON  553 

carbon  monoxide,  which  then  reduces  the  ferric  oxide  to  metallic 
iron  according  to  the  equation  : 

Fe203  4-  3CO->2Fe  +  3CO2. 

The  reduction  of  the  oxide  takes  place  mainly  in  the  upper  part  of 
the  furnace,  where  the  temperature  is  not  sufficiently  high  (600°  to 
900°)  to  fuse  the  iron.  On  its  way  down  the  furnace,  the  finely 
divided  iron  reaches  the  regions  of  higher  temperature  at  ^,  where 
it  begins  to  take  up  carbon,  whereby  its  melting-point  is  lowered.1 
Finally  it  becomes  completely  fused  and  collects  in  the  cavity  h, 
at  the  bottom  of  the  furnace.  At  the  same  stage  (at  <?),  the  clay  and 
silica  originally  present  in  the  ore  combine  with  the  lime  to 
form  a  fusible  slag  (a  mixture  of  calcium  and  aluminium  silicates) 
which,  being  lighter  than  the  iron,  forms  a  layer  above  it.  The 
iron  is  run  off  periodically  into  moulds,  forming  the  so-called  pig- 
iron;  the  slag  runs  off  continually  through  an  opening  at  a  higher 
level.  A  blast-furnace  can  be  worked  for  years  without  stopping, 
being  kept  nearly  full  by  admission  of  fresh  charges  as  required. 

All  the  iron  used  in  commerce  contains  other  substances  in  greater 
or  less  proportion,  which  greatly  modify  its  properties.  The  more 
important  varieties  of  iron  in  common  use  are  cast  iron^  wrought  iron 
and  steel.  Wrought  iron  is  nearly  pure  iron. 

Cast  Iron — The  most  important  alloys  of  iron  are  those  con- 
taining carbon.  The  carbon  may  be  present  in  three  different  forms 
(a)  as  graphite  ;  (b)  as  a  definite  carbide  of  iron,  Fe3C  ;  (3)  in  a 
state  of  solid  solution  in  the  iron.  In  the  fused  state  iron  dissolves 
a  considerable  proportion  of  carbon  ;  a  solution  saturated  at  1400° 
contains  about  6  per  cent,  of  dissolved  carbon.  When  iron  containing 
carbon  in  solution  is  cooled  slowly,  the  greater  part  of  the  carbon 
separates  as  graphite,  but  when  it  is  cooled  rapidly,  the  carbon  re- 
mains partly  as  carbide  and  is  partly  retained  in  solid  solution. 

The  alloy  containing  admixed  graphite  is  known  as  grey  cast  iron; 
when  it  is  treated  with  dilute  acids  the  graphite  remains  as  a  residue. 
The  alloy  in  which  all  the  carbon  is  combined  or  dissolved  is  known 
as  white  cast  iron^  being  silvery  white  in  colour  ;  when  treated  with 
acids  the  carbide  is  decomposed  and  hydrocarbon  gases,  with  a 
disagreeable  odour,  are  given  off.  Grey  cast  iron  is  soft  and  gives 
good  castings;  it  is  also  used  in  the  preparation  of  steel 
WThite  cast  iron  is  very  hard  and  brittle. 

1  Pure  iron  melts  at  1500°  ;  iron  containing  4.3  per  cent,  of  carbon  at  1125°. 


554     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Besides  carbon,  cast  iron  contains  silicon,  phosphorus,  sulphur, 
manganese  and  other  elements  in  smaller  proportion.  These  ele- 
ments are  reduced  from  their  oxygen  compounds  by  carbon  in  the 
hottest  part  of  the  furnace,  and  are  then  taken  up  by  the  iron. 

Cast  iron  is  obtained  by  melting  pig-iron  and  casting  it  in  moulds  ; 
its  properties  depend  on  the  conditions  under  which  the  melting  and 
casting  are  performed. 

Wrought  Iron — This  is  nearly  pure  iron,  and  is  obtained  from 
cast  iron  by  removing  the  impurities  by  oxidation.  The  process  used 
is  termed  puddling.  Cast  iron  is  melted  in  a  reverberatory  furnace 
(p.  404),  the  hearth  of  which  is  lined  with  ferric  oxide.  The  mass 
is  continuously  stirred  or  puddled,  in  order  to  secure  thorough  mixing 
with  the  ferric  oxide.  In  the  process  carbon  is  oxidized  to  carbon 
dioxide,  which  escapes  ;  the  silicon  and  phosphorus,  after  oxidation, 
unite  with  the  ferric  oxide  to  form  a  slag.  Meanwhile,  owing  to  the 
removal  of  the  impurities,  the  iron  has  become  pasty  ;  it  is  then 
rolled  up  into  large  balls,  which  are  removed  from  the  furnace  and 
placed  under  a  steam  hammer,  by  means  of  which  the  slag  is  squeezed 
out  and  the  particles  of  metal  formed  into  a  solid  mass. 

Wrought  iron  contains  up  to  0.2  per  cent,  of  carbon,  and  traces  of 
silicon  and  phosphorus.  Although  it  fuses  only  with  great  difficulty 
(pure  iron  melts  about  1500°)  it  can  be  softened  by  heat  at  a  much 
lower  temperature  and  two  pieces  joined  by  hammering — a  process 
known  as  welding. 

Steel — Steel  contains  0.3  to  rather  over  1.5  per  cent,  of  carbon, 
according  to  the  purpose  for  which  it  is  to  be  used,  and  is  practically 
free  from  silicon,  phosphorus  and  sulphur.  It  is  prepared  from  cast 
iron  by  removing  the  carbon,  silicon,  etc.,  by  oxidation.  Two  prin- 
cipal methods  for  this  purpose  are  in  use  (i)  the  Bessemer  process ; 
(2)  the  Siemens- Mar  tin  or  open-hearth  process. 

(i)  The  Bessemer  process  is  carried  out  in  a  converter  of  the  type 
shown  in  Fig.  90.  The  converter  is  made  of  iron  plates,  and  is 
suspended  on  trunnions  ;  at  the  bottom  there  are  a  number  of  small 
holes  through  which  air,  supplied  by  the  pipe  shown  in  the  diagram, 
can  be  blown.  The  converter  is  tilted  into  a  horizontal  position,  the 
blast  is  started,  a  charge  of  molten  iron  (about  10  tons)  is  run  in, 
and  the  converter  tilted  back  into  a  vertical  position.  The  carbon  is 
converted  to  carbon  monoxide,  which  burns  with  a  long  flame  at 
the  mouth  of  the  converter ;  the  other  impurities  are  oxidized  and 
combine  to  form  a  slag,  part  of  the  lining  being  used  up  in  this 
reaction.  The  end  of  the  operation  is  indicated  by  the  disappear- 


IRON 


555 


ance  of  the  carbon  monoxide  flame.  Sufficient  ferro-manganese  or 
spiegel-eisen  (an  alloy  containing  iron,  manganese,  and  a  fairly  high 
proportion  of  carbon)  is  then  added  to  deoxidize  the  iron  and  to  obtain 
the  proper  proportion  of  carbon,  and  the  steel  poured  into  moulds. 

In  the  original  Bessemer  process,  the  converter  was  lined  with 
silica,  but  it  was  found  that  under  these  circumstances  phosphorus, 
the  presence  of  which  in  steel  is  very  deleterious,  was  not'  removed. 
This  difficulty  is  overcome  by  the  use  of  the  so-called  basic  lining, 
a  mixture  of  calcium  and  magnesium  oxides,  as  suggested  by  Thomas 
and  Gilchrist.  The  phosphorus  combines  with  the  lining  to  form 


FIG.  90. 

calcium  and  magnesium  phosphates,  which  pass  into  the  slag,  con- 
stituting the  so-called  basic  slag,  which  is  largely  used  as  a  manure. 

(2)  In  the  Siemens- Mar  tin  process,  the  pig  iron,  along  with  some 
scrap  iron,  is  melted  on  the  hearth  of  a  furnace,  an  oxide  of  iron, 
generally  haematite,  is  added,  and  a  flame  of  producer  gas  (p.  333)  is 
caused  to  play  across  the  surface  for  some  hours.  The  oxidation  ol 
impurities  is  effected  partly  by  the  oxygen  of  the  haematite,  partly  by 
that  in  the  hot  flame.  In  this  case  also,  the  required  proportion  of 
carbon  is  added  in  the  form  of  spiegel-eisen. 

Steel  is  also  prepared  from  wrought  iron  by  the  so-called  cementation 
process.  The  bars  of  iron  are  imbedded  in  charcoal  and  exposed  for 
several  days  to  a  high  temperature  in  large  chests  of  firebrick.  The 


556     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


mode  in  which  the  carbon  is  conveyed  into  the  interior  of  the  iron 
bars  in  the  cementation  process  is  not  understood. 

The  most  important  property  of  steel  is  that  when  heated  and 
suddenly  cooled  it  becomes  hard  and  brittle,  when  it  is  said  to  be 
hardened.  Pure  iron  does  not  become  appreciably  harder  under 
these  conditions.  The  hardness  of  steel  may  be  reduced  to  any 
desired  extent  by  raising  to  a  moderate  temperature  and  then  allow- 
ing to  cool — a  process  known  as  tempering.  The  temper  depends 
upon  the  temperature  to  which  the  steel  is  raised  (the  higher  the 
tempering  temperature  the  softer  the  steel),  and  this  is  judged  by 
the  operator  according  to  the  tint  which  the  surface  of  the  metal 
assumes  during  the  operation.  A  light  straw  colour  (420-450°)  serves 
for  tempering  razors,  a  purple  colour  (500°)  for  cutlery,  and  a  dark 
blue  (600°)  for  chisels  and  saws. 

The  properties  of  steel  also  depend  upon  the  proportion  ot  carbon 
it  contains.  A  mild  steel  contains  0.25  to  0.5  per  cent,  of  carbon,  the 
hardest  steel  (tool  steel)  about  I  per  cent,  of  carbon.  In  steel  the 
carbon  is  all  present  in  chemical  combination  or  in  solution.  The 
average  composition  of  typical  specimens  of  wrought  iron,  grey  cast 
iron,  white  cast  iron  and  steel  is  given  in  the  accompanying  table  : — 


Cast 

Cast 

Wrought 

Steel 

Steel 

(Grey). 

(White). 

Iron. 

(Mild). 

(Hard). 

Carbon  (graphite)    . 
„       (combined)  . 

3-3 
0.4 

4.1 

|a,4{ 

0.21 

1.20 

Silicon             .         '. 

I  Q 

O.2 

0.  14 

o  06 

o.  i  ^ 

Sulphur                t     . 

O.O2 

O.O^ 

O.O2 

o.os 

O.OI 

Phosphorus    .     .     . 

0.6 

0.07 

O.IO 

0.02 

0.02 

Manganese     .     .     . 

0.4 

2.0 

0.15 

0.52 

0.25 

The  changes  taking  place  in  iron  and  steel  under  different  condi- 
tions of  cooling  are  very  complicated  and  by  no  means  thoroughly 
understood  ;  they  can  only  be  dealt  with  very  briefly  here.  Iron  is 
considered  to  exist  in  three  allotropic  modifications  :  y-iron,  which  is 
stable  above  900° ;  $-iron,  stable  between  900°  and  760°  ;  and  a-iron, 
stable  below  760°.  y-iron  and  /3-iron  are  non-magnetic,  y-iron  can 
retain  a  considerable  proportion  of  carbon  in  solid  solution  ;  /?  and 
a-iron  dissolve  little  or  no  carbon.  When  y-iron  containing  dissolved 
carbon  is  very  rapidly  cooled  it  forms  a  hard,  brittle  alloy  (hard  ened 
steel) ;  it  is  highly  supercooled,  and  the  presence  of  carbon  retards 


IRON  557 

its  transformation  to  the  modification  (a-iron)  stable  under  these  condi- 
tions. When,  on  the  other  hand,  the  iron  containing  carbon  in  solution 
is  cooled  so  slowly  that  equilibrium  is  attained  at  every  stage,  the  final 
product  contains  a-iron,  a  little  carbide,  Fe3C  (known  as  cementite],  and 
graphite.  The  moderate  heating  and  subsequent  cooling  used  in  the 
tempering  of  steel  brings  the  alloy  towards  the  equilibrium  state 
to  an  extent  determined  by  the  conditions.  The  properties  of  a  steel 
depend  upon  the  relative  proportions  of  -y-iron,  solid  solution  of  carbon 
in  iron,  cementite,  and  a  and  /3-iron  in  the  product,  and  on  their 
state  of  distribution.  Pearlite^  a  name  often  met  with  in  the  discus- 
sion of  steels,  denotes  a  eutectic  of  cementite  and  a-iron,  and  contains 
0.9  per  cent,  of  carbon. 

Properties  of  Pure  Iron  —  Pure  iron  can  be  obtained  by 
heating  ferric  oxide  in  hydrogen,  or  by  electrolysis  of  an  iron  salt, 
e.g.  ferric  chloride,  in  aqueous  solution.  It  is  a  white  lustrous  metal 
which  takes  a  high  polish  ;  its  density  is  7.86.  It  is  malleable  and 
ductile,  and  is  not  very  hard.  It  melts  about  1500°,  but  softens  at 
much  lower  temperatures.  It  is  magnetic,  but,  unlike  steel,  it  rapidly 
loses  this  property  when  the  magnetizing  force  is  withdrawn. 

Iron  is  permanent  in  dry  air,  but  in  moist  air  it  quickly  becomes 
coated  with  a  layer  of  ferrous  and  ferric  oxides,  and  is  said  to  rust. 
As  the  rust  does  not  form  a  continuous  coating,  but  peels  off  in  sheets, 
the  corrosion  spreads  to  deeper  layers  of  the  metal.  Rusting  is 
greatly  accelerated  by  the  presence  of  carbon  dioxide  (from  the  air). 

The  exact  mechanism  of  the  rusting  of  iron  is  at  present  a  matter 
of  dispute.  Grace  Calvert  and,  later,  Crum  Brown  expressed  the 
view  that  carbon  dioxide  is  essential;  the  first  stage  of  the  process  is 
the  formation  of  ferrous  carbonate,  FeCO3;  thus  — 


which  is  then  oxidized  more  or  less  completely  to  red  ferric  oxide. 
Moody  (Trans.  Chem.  Soc.,  1906,  89,  720)  adopts  this  view,  and  con- 
tends that  iron  does  not  rust  in  air  and  water  entirely  freed  from 
carbon  dioxide.  In  the  presence  of  carbon  dioxide  and  water  the 
iron  dissolves  to  form  ferrous  bicarbonate,  which  then  undergoes 
oxidation  as  stated  above.  Lambert  and  Thomson  (Trans.  Cfiem. 
Soc.,  1910,  97,  2426),  on  the  other  hand,  state  that  although  rusting 
does  not  occur  with  perfectly  pure  iron  in  contact  with  pure  water 
and  pure  oxygen,  traces  of  impurities  are  sufficient  to  cause  oxidation 
under  the  same  conditions,  even  if  the  impurity  be  not  of  an  acid 
nature  or  likely  to  produce  an  acid  during  the  reaction.  Walker, 


558     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

Cederholm  and  Bent  (Trans.  Amer.  Chem.  Soc.,  1907,  xxix.  1251)  also 
state  that  carbon  dioxide  is  not  essential  for  rusting  to  occur,  and  advo- 
cate an  electrolytic  explanation  of  the  phenomenon.  In  virtue  of  its 
solution  pressure  (p.  85)  iron  sends  out  Fe"  ions  even  into  pure  water  ; 
but  in  the  absence  of  oxygen  the  reaction  soon  comes  to  an  end,  owing 
to  the  accumulation  of  gaseous  hydrogen  on  the  metal,  which  sets  up 
an  E.M.F.  of  polarization.  In  the  presence  of  oxygen,  however, 
the  hydrogen  is  removed  by  oxidation,  and  the  ferrous  salt  is  removed 
from  the  solution  by  precipitation  as  oxide,  so  that  the  reaction  is 
enabled  to  proceed. 

Iron  readily  dissolves  in  dilute  hydrochloric  and  sulphuric  acids, 
with  evolution  of  hydrogen  and  formation  of  the  corresponding  salts. 
With  dilute  nitric  acid,  ferrous  nitrate  and  ammonium  nitrate  are 
formed.  Iron  is  not  attacked  by  concentrated  nitric  acid,  but  is 
changed  to  the  passive  state,  in  which  condition  it  no  longer  precipi- 
tates copper  from  its  salts.  This  passive  condition  is  also  brought 
about  by  other  oxidizing  agents,  e.g.  potassium  dichromate,  hydrogen 
peroxide.  The  explanation  of  the  passive  state  most  in  favour  is  that 
it  is  due  to  the  presence  of  a  thin  film  of  oxide  on  the  surface  of  the 
metal.  There  is  a  considerable  difference  of  potential  between  a 
metal  in  the  active  and  the  same  metal  in  the  passive  state. 

Iron  Alloys — The  alloys  with  carbon  have  already  been  fully 
dealt  with.  Certain  of  the  properties  of  steel  can  be  enhanced  by  the 
addition  of  other  elements.  The  addition  of  manganese  (about  12  per 
cent.)  confers  increased  hardness  on  steel,  while  the  alloy  possesses  a 
high  degree  of  ductility  and  elasticity.  Nickel  (up  to  3  per  cent.) 
makes  steel  more  tough  and  elastic.  Chrome  steel  (containing  about 
2  per  cent,  of  chromium)  combines  intense  hardness  with  a  high 
elastic  limit.  Tungsten  is  also  a  useful  addition  under  certain 
circumstances. 


OXIDES  AND  HYDROXIDES  OF  IRON 

i 

Three  oxides  of  iron  are  known  : 

Ferrous  oxide FeO 

Ferric  oxide  .......     Fe2O3 

Triferric  tetroxide  (magnetic  oxide  of  iron)  .     Fe3O4 

The  first  two  oxides  are  basic,  giving  rise  to  ferrous  salts  (type 
FeCl2)  and  ferric  salts  (type  FeCl3)  respectively.  The  third  is  a  mixed 
oxide,  yielding  a  mixture  of  ferrous  and  ferric  oxides,  FeO,  Fe2O3. 


IRON  559 

Salts  derived  from  an  acidic  oxide,  FeO3,  are  also  known,  e.g. 
potassium  ferrate,  K2FeO4. 

Ferrous  Oxide,  FeO,  is  obtained  by  heating  ferric  oxide  in 
hydrogen  at  300°,  or  by  heating  ferrous  oxalate  out  of  contact  with  air. 
It  is  a  black  powder,  which  readily  becomes  oxidized  on  exposure  to 
air,  and  dissolves  in  acids  to  form  ferrous  salts. 

Ferrous  Hydroxide,  Fe(OH)2,  is  obtained  as  a  white  precipi- 
tate by  mixing  air-free  solutions  of  a  ferrous  salt  and  an  alkali 
hydroxide  in  the  absence  of  air.  It  rapidly  turns  green  in  the  air 
owing  to  oxidation,  and  finally  becomes  brown  owing  to  the  formation 
of  ferric  hydroxide. 

Ferric  Oxide,  Fe2O3,  occurs  in  lustrous  black,  six-sided  crystals, 
as  specular  iron  ore,  and  in  reddish  masses  as  hcematite.  It  is  obtained 
as  a  red  amorphous  powder  by  heating  ferrous  sulphate  in  the  prepa- 
ration of  Nordhausen  sulphuric  acid,  and  is  also  obtained  when  most 
iron  salts  are  strongly  heated  in  air.  It  is  stable  in  air  at  a  red  heat, 
but  above  1000°  the  compound  Fe3O4  is  formed.  It  dissolves  in  acids 
(though  very  slowly,  especially  after  strong  ignition)  to  form  ferric 
salts.  The  amorphous  powder  is  used  as  jewellers'  rouge  for  polish- 
ing purposes,  and  also  as  a  pigment  under  the  name  of  Venetian  red. 

Ferric  Hydroxide,  Fe(OH)3vrH2O,  is  obtained  as  a  brown 
gelatinous  precipitate  by  adding  excess  of  ammonium  hydroxide  to  a 
solution  of  a  ferric  salt.  By  careful  dehydration  a  compound  of  the 
composition  Fe(OH)3or  Fe2O3,3H2O  appears  to  be  obtained.  The 
water  in  this,  as  in  other  cases  is  adsorbed  by  the  hydroxide,  and  it 
is  doubtful  whether  definite  compounds  are  formed  (cf.  p.  482).  The 
compound  2Fe2O3,3H2O  occurs  in  nature  as  limonite  (brown  hcematite). 
Iron  rust  is  mainly  a  hydrated  ferric  oxide,  with  ferrous  oxide  and 
carbonate  in  smaller  proportion. 

Freshly  precipitated  and  washed  ferric  hydroxide  is  soluble  in  a 
solution  of  ferric  chloride.  The  iatter  can  be  almost  completely 
removed  by  prolonged  dialysis  (p.  369),  and  a  deep  brown,  practically 
tasteless  colloidal  solution  of  ferric  hydroxide,  the  so-called  dialysed 
iron,  is  obtained. 

Ferric  hydroxide  is  a  very  weak  base,  and  therefore  the  ferric  salts 
are  considerably  hydrolyzed  in  aqueous  solution. 

Triferric  Tetroxide  {magnetic  iron  ore),  Fe3O4,  occurs  natu- 
rally in  nearly  black  octahedral  crystals,  as  magnetite  or  lodestone, 
so  called  because  it  is  magnetic.  It  is  formed  when  iron  is  burned 
at  a  high  temperature  in  oxygen  or  air  (hammer  scale),  and  also  when 
steam  is  passed  over  heated  iron  (p.  34).  When  heated  with  acids  it 


560     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

yields  a  mixture  of  ferrous  and  ferric  salts,  and  may  therefore  be 
represented  as  FeO,Fe2O3  (cf.  chrome  iron  ore,  FeO,Cr2O3). 

Ferrates  —  When  chlorine  is  passed  into  ferric  hydroxide  sus- 
pended in  cold  concentrated  potassium  hydroxide  solution,  a  dark  red 
solution  is  obtained,  from  which,  on  cautious  evaporation,  potassium 
ferrate,  K2FeO4,  can  be  obtained  in  dark  red  crystals,  isomorphous 
with  potassium  sulphate  and  potassium  chromate  : 


Potassium  ferrate  is  derived  from,  ferric  acid,  H2FeO4,  but  neither  the 
acid  nor  the  corresponding  oxide,  FeO3  (cf.  SO3,  CrO3),  has  been  iso- 
lated. When  the  solution  of  a  ferrate  is  acidified-,  oxygen  is  given 
off  and  ferric  and  potassium  salts  remain  in  solution. 

FERROUS  SALTS 

Most  ferrous  salts  are  light  green  in  colour.  They  are  obtained  by 
dissolving  the  metal  in  the  appropriate  acid  or  by  reducing  ferric  salts 
in  acid  solution  by  "  nascent  "  hydrogen,  hydrogen  sulphide,  etc. 
They  are  easily  oxidized  to  ferric  salts. 

Ferrous  Chloride,  FeCl2,  is  obtained  in  the  anhydrous  form  in 
colourless  crystals  by  heating  iron  in  a  current  of  dry  hydrogen 
chloride.  It  is  obtained  as  FeCl2,4.H2O  in  green,  monoclinic  crystals, 
by  dissolving  iron  in  hydrochloric  acid,  and  evaporating  the  solution 
with  exclusion  of  air.  When  it  is  heated  in  the  air,  a  mixture  of 
ferric  oxide  and  chloride  is  obtained  : 

6FeCl2+3O-^4FeCl3+Fe2O3. 

Ferrous  Sulphate  (Green  Vitriol],  FeSO4,7H2O,  is  obtained  by 
dissolving  iron  in  dilute  sulphuric  acid.  Commercially  it  is  prepared 
by  exposing  iron  pyrites,  FeS2,  to  atmospheric  oxidation  ;  the  result- 
ing ferrous  sulphate  is  extracted  with  water  and  the  solution  evapo- 
rated. It  separates  from  solution  in  green,  monoclinic  prisms, 
isomorphous  with  ZnSO4,7H2O  and  MgSO4,7H2O.  When  heated  to 
100°  the  crystals  lose  6H2O  ;  at  a  higher  temperature  the  monohy- 
drate,  FeSO4,H2O,  is  decomposed  and  ferric  oxide  is  obtained. 

When  exposed  to  air  at  room  temperature,  the  heptahydrate  efflor- 
esces slightly  and  becomes  coated  with  a  brown  layer  of  basic  ferric 
sulphate,  Fe2O(SO4)2  (or  perhaps  Fe(OH)SO4).  The  same  product 
is  formed  when  the  heptahydrate  is  roasted  in  the  air.  The  basic 
sulphate  obtained  by  roasting  was  formerly  used  for  preparing  Nord- 
hausen  sulphuric  acid  ;  for  this  purpose  it  was  distilled  from  clay 


IRON  561 

retorts  and  the  volatile  products  collected  in  water  or  sulphuric  acid 
(p.  302)  : 

Fe2O(SO4)2->Fe2O3 


Ferrous  sulphate  forms  double  salts  with  the  alkali  sulphates.  The 
most  important  is  ferrous  ammonium  sulphate,FeSO4,(NH4)2SO4,6H2O, 
which  is  less  easily  oxidized  in  the  air  than  ferrous  sulphate,  and 
therefore  finds  application  in  volumetric  analysis. 

Ferrous  sulphate  is  used  as  a  disinfectant,  in  the  preparation  of 
ink,  etc. 

Ferrous  Carbonate,  FeCO3,  occurs  naturally  as  spathic  iron  ore> 
in  rhombohedral  crystals  isomorphous  with  calc-spar.  It  is  obtained 
by  double  decomposition  when  Fe"  and  CO3"  ions  are  brought  to- 
gether in  solution,  but  is  very  rapidly  oxidized  in  the  air.  It  is  used 
in  medicine  mixed  with  sugar,  which  to  some  extent  protects  it  against 
oxidation. 

Ferric  carbonate  has  not  been  obtained.  When  Fe"f  and  CO3" 
ions  are  brought  together  in  solution,  the  ferric  carbonate  which  may 
be  momentarily  formed  is  completely  hydrolyzed  by  water  and  ferric 
hydroxide  is  precipitated. 

FERRIC  SALTS 

Ferric  Hydroxide,  Fe(OH)3,  is  a  very  weak  base,  and  therefore 
ferric  salts  with  strong  acids  are  considerably  hydrolyzed  in  solution, 
and  salts  with  weak  acids  (e.g.  carbonic  acid,  hydrogen  sulphide) 
cannot  be  obtained  from  aqueous  solution. 

Ferric  Chloride,  FeCl3,  is  obtained  in  the  anhydrous  form,  as 
lustrous,  dark  green  crystals,  by  heating  iron  in  a  current  of  chlorine 
gas.  It  is  obtained  in  solution  by  oxidizing  ferrous  chloride  with 
chlorine,  nitric  acid  or  other  oxidizing  agent,  and  separates  from 
solution  as  the  hexahydrate,  FeCl3,6H2O,  in  yellow  crystals.  Lower 
hydrates  of  the  salt  are  also  known^  It  cannot  be  obtained  in  the 
anhydrous  form  by  heating  one  of  the  hydrates  in  air,  as  at  high 
temperatures  hydrochloric  acid  is  given  off. 

The  anhydrous  salt  is  fairly  volatile.  At  300  to  400°  the  vapour 
density  corresponds  approximately  with  the  formula  Fe2Cl6 ;  above 
750°  with  the  formula  FeCl3. 

Ferric  chloride  is  considerably  hydrolyzed  in  solution,  the  hydroxide 
remaining  dissolved  in  colloidal  form  (hydrosol)  : 


562     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Ferric  Sulphate,  Fe2(SO4)3,  is  obtained  by  oxidizing  ferrous 
sulphate  in  sulphuric  acid  solution.  Nitric  acid  may  conveniently  be 
used  for  this  purpose  : 


Adding 


The  aqueous  solution  is  brownish-red  (owing  to  the  presence  of 
colloidal  ferric  hydroxide,  formed  by  hydrolysis)  ;  on  evaporation 
the  salt  separates  in  the  anhydrous  form  as  a  grayish-white  powder. 

When  the  calculated  quantity  of  potassium  or  ammonium  sulphate 
is  added  to  ferric  sulphate  solution,  and  the  solution  is  evaporated 
over  sulphuric  acid  at  room  temperature,  an  iron  alum,  e.g. 
Fe2(SO4)3,K2SO4,24H2O,  separates  in  violet,  octahedral  crystals 
(p.  484). 

SULPHIDES  OF  IRON 

Ferrous  Sulphide,  FeS,  is  obtained  as  a  dark  metallic-looking 
mass  by  heating  iron  and  sulphur  together.  The  sulphide  thus  ob- 
tained is  dissolved  by  acids,  hydrogen  sulphide  being  given  off  and 
a  ferrous  salt  formed.  Hydrogen  sulphide  is  usually  prepared  by 
this  method. 

It  follows  from  the  above  that  no  precipitate  is  obtained  when 
hydrogen  sulphide  is  passed  into  the  solution  of  a  ferrous  salt  con- 
taining free  acid,  but  the  sulphide  is  obtained  as  a  black  amorphous 
precipitate  when  ammonium  sulphide  is  added  to  a  ferrous  or  ferric 
salt.  When  a  ferric  salt  is  used,  free  sulphur  is  also  formed  : 

2FeCl3+3(NH4)2S->6NH4Cl  +  2FeS  +  S. 

The  precipitation  is  preceded  by  the  reduction  of  the  iron  from  the 
ferric  to  the  ferrous  condition.  The  amorphous  sulphide  when  moist 
is  rapidly  oxidized  in  the  air  to  ferrous  sulphate. 

Ferric  Sulphide,  Fe2S3,  cannot  be  obtained  by  precipitation  in 
presence  of  water  (see  above),  but  is  said  to  be  formed  as  a  yellow 
mass  when  a  mixture  of  the  components  in  the  calculated  proportions 
is  heated. 

Iron  Disulphide,  FeS2,  occurs  naturally  as  iron  pyrites  in  yellow 
crystals  belonging  to  the  regular  system  (cubes,  octahedra  or  other 
forms),  and  is  also  formed  by  gently  heating  a  mixture  of  ferrous 


IRON  563 

sulphide  and  sulphur.  When  heated  in  air  it  burns  to  sulphur  dioxide 
and  ferric  oxide  : 

4FeS2+iiO2->2Fe2O3  +  8SO2, 

and  is  largely  used  in  preparing  sulphuric  acid  (p.  307).  The  other 
sulphides  of  iron  also  give  ferric  oxide  and  sulphur  dioxide  when 
roasted  in  air. 

When  it  is  heated  in  absence  of  air,  part  of  the  sulphur  is  given  off 
and  ferroso-ferric  sulphide,  Fe3S4,  is  formed: 


COMPLEX  IRON  CYANOGEN  COMPOUNDS 

The  normal  cyanides  of  iron,  Fe(CN)2  and  Fe(CN)3,  are  not 
known,  but  two  complex  cyanides,  potassium  ferro  cyanide,  K4Fe(CN)6, 
and  potassium  ferricyanide,  K3Fe(CN)6,  are  known.  Potassium 
ferrocyanide  is  obtained  by  fusing  together  potassium  carbonate,  iron 
filings  and  animal  refuse  such  as  charred  blood,  scraps  of  horns, 
hoofs,  etc.  ;  the  mass  is  extracted  with  water,  and  on  evaporation 
potassium  ferrocyanide  separates  in  the  form  of  light  yellow  crystals 
as  K4Fe(CN)6,3H2O.  When  potassium  ferrocyanide  is  strongly 
heated,  potassium  cyanide,  nitrogen  and  iron  carbide  are  formed  : 

K4Fe(CN)6->4KCN  +  FeC2+  N2. 

Its  employment  in  the  preparation  of  carbon  monoxide  and  of  hydro- 
cyanic acid  has  already  been  mentioned. 

Potassium  Ferricyanide,  K3Fe(CN)6,  is  obtained  bypassing 
chlorine  into  a  solution  of  potassium  ferrocyanide  : 

2K4Fe(C  N  )6  +  Cl2->2K3Fe(C  N)6  +  2  KC1. 

The  ferricyanide  occurs  in  dark-red  crystals  which  are  readily  soluble 
in  water  ;  it  acts  as  an  oxidizing  agent. 

It  might  appear  that  these  salts  can  be  regarded  as  compounds  ot 
potassium  cyanide  with  ferrous  and  ferric  cyanide  respectively,  thus  — 
K4Fe(CN)6=4KCN,Fe(CN)2  and  K3Fe(CN)6  =  3KCN,Fe(CN)3.  As 
a  matter  of  fact,  however,  the  solutions  contain  the  complex  ions 
Fe(CN)0////  and  Fe(CN)6///  respectively  ;  in  other  words  the  group 
Fe(CN)6has  four  negative  charges  in  the  ferro-  and  three  in  the  ferri- 
cyanides.  The  solutions  give  none  of  the  reactions  for  iron  salts  (so 
that  Fe"  and  Fe*1'  ions  are  absent)  and  are  not  poisonous  (absence 
ofCN'ion). 


564     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Potassium  ferrocyanide  and  ferricyanide  are  used  to  distinguish 
between  ferrous  and  ferric  salts.  Potassium  ferrocyanide  gives 
with  a  ferrous  salt  a  light  blue  precipitate  of  ferrous  ferrocyanide, 
Fe2"Fe(CN)6"" ;  with  a  ferric  salt  a  dark  blue  precipitate  of  ferric 
ferrocyanide,  Fe4"'3[Fe(CN)6]"",  known  as  Prussian  blue.  Potassium 
ferricyanide  gives  with  ferrous  salts  a  dark  blue  precipitate  of  ferrous 
ferricyanide,  Fe3"2[Fe(CN)6]'";  with  ferric  salts  a  dark  red  colour, 
but  no  precipitate  is  obtained.  The  formulae  for  these  compounds 
can  readily  be  understood  when  it  is  noted  that  the  ferrocyanides 
are  derived  from  the  tetrabasic  hydroferrocyanic  add,  H4Fe(CN)6, 
the  ferricyanides  from  the  tribasic  hydroferricyanic  acid,  H3Fe(CN)6. 
The  blue  compounds  are  decomposed  by  alkalis  into  the  correspond- 
ing hydroxide  and  alkali  ferro-  or  ferricyanide,  the  colour  being 
destroyed. 

Tests  for  Iron — The  majority  of  the  tests  for  iron  have  been 
mentioned  in  the  course  of  the  chapter.  Iron  salts,  whether  ferrous 
or  ferric,  give  with  ammonium  sulphide  a  black  precipitate  of  ferrous 
sulphide,  soluble  in  dilute  acids.  The  behaviour  of  ferrous  and  ferric 
salts  with  ammonia  and  with  potassium  ferrocyanide  and  ferricyanide 
is  characteristic.  Ferric  salts  give  with  an  alkali  thiocyanate  a  deep 
red  coloration,  due  to  ferric  thiocyanate,  Fe(CNS)3. 


COBALT 

Symbol,  Co.     Atomic  weight  =  58. 97. 

Occurrence — Cobalt  is  found  free,  along  with  iron,  nickel,  and 
other  metals,  in  meteorites.  In  the  combined  state,  it  occurs  mainly 
as  smaltite,  CoAs2,  and  cobalt  glance,  CoAsS. 

Preparation  of  Metal — The  metal  is  obtained  by  reducing 
one  of  the  oxides  in  a  current  of  hydrogen,  or  better,  by  reducing  the 
oxide  with  aluminium  powder  (Goldschmidt's  process). 

Properties — Cobalt  is  a  white,  lustrous,  tenacious  metal  capable 
of  taking  a  high  polish  ;  its  density  is  8.5.  It  melts  at  1460°,  about 
40°  below  iron.  It  is  malleable,  and  when  heated  is  ductile.  It  is 
magnetic,  but  not  so  strongly  so  as  iron.  It  is  practically  unaffected 
by  air  at  room  temperature,  but  on  heating  strongly  forms  the  oxide 
Co3O4.  Hydrochloric  and  sulphuric  acids  dissolve  it  very  slowly, 
but  it  is  rapidly  dissolved  by  nitric  acid,  with  formation  of  cobaltous 
nitrate,  Co(NO3)2. 


COBALT  565 


OXIDES  AND  HYDROXIDES  OF  COBALT 

Three  oxides  of  cobalt  are  definitely  known  :  (a)  Cobaltous  oxide, 
CoO  ;  (b)  cobaltic  oxide,  Co2O3  ;  and  (c)  cobalto-cobaltic  oxide,  Co3O^ 
Cobalt  dioxide,  CoO2,  has  also  been  described,  but  does  not  appear 
to  have  been  obtained  in  a  pure  condition.  Like  iron,  cobalt  forms 
two  series  of  salts,  cobaltous  salts,  in  which  it  is  divalent,  and  cobaltic 
salts,  in  which  it  is  trivalent.  The  cobaltic  salts  are,  however, 
extremely  unstable. 

Cobaltous  Oxide  (cobalt  monoxide),  CoO,  is  obtained  as  a 
green  powder  by  heating  the  hydroxide  or  carbonate  in  absence  of 
air.  It  is  stable  in  the  air  at  room  temperature,  but  at  a  red  heat 
takes  up  oxygen  and  forms  cobalto-cobaltic  oxide,  Co3O4. 

Cobaltous  Hydroxide,  Co(OH)2,  is  obtained  as  a  red  pre- 
cipitate by  adding  an  alkali  hydroxide  to  the  solution  of  a  cobaltous 
salt  and  then  heating  to  decompose  the  basic  salt  first  obtained. 
The  hydroxide  soon  turns  brown  in  the  air  owing  to  oxidation.  It  is 
insoluble  in  excess  of  alkali  hydroxide,  but  dissolves  in  ammonia, 
with  formation  of  complex  compounds  (see  cobaltammine  compounds, 
below). 

Cobaltic  Oxide,  Co2O3,  is  obtained  as  a  black  powder  by 
cautiously  heating  cobaltous  nitrate.  At  a  red  heat  it  loses  a  little 
oxygen,  forming  the  compound  Co3O4.  Cobaltic  oxide  dissolves  in 
sulphuric  or  hydrochloric  acid  in  the  cold,  and  the  solutions  pre- 
sumably contain  unstable  cobaltic  salts.  On  heating  the  sulphate 
solution  oxygen  is  given  off,  whilst  under  the  same  circumstances  the 
chloride  solution  gives  off  chlorine  ;  and  in  both  cases  cobaltous  salts 
remain  in  solution  (cf.  manganese  dioxide,  p.  544). 

Cobaltic  Hydroxide,  Co(OH)3,  is  formed  as  a  dark  pre- 
cipitate by  the  action  of  sodium  hypochlorite  on  cobaltous  hydroxide. 

Cobalto-Cobaltic  Oxide,  Co3O4,  the  oxide  usually  met  with 
in  commerce,  occurs  as  a  black  powder.  With  acids  it  behaves  like 
a  mixture  of  cobaltous  and  cobaltic  oxides. 

COBALTOUS  SALTS 

The  cobaltous  salts  are  the  only  stable  salts  of  cobalt.  The  solid 
hydrates  are  pink,  and  they  form  pink  solutions  at  room  temperature  ; 
the  anhydrous  salts  are  generally  blue.  The  cause  of  the  differences 
in  colour  are  doubtless  due  to  differences  in  constitution  referred  to 
later  in  connexion  with  the  complex  ammines  (p.  567). 


566     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Cobaltous  Chloride,  CoCl2,  is  obtained  by  dissolving  any  of 
the  oxides  of  cobalt  in  hydrochloric  acid,  and  separates  on  evapo- 
rating the  solution  in  red  monoclinic  crystals  as  CoCl2,6H2O.  On 
heating  to  120°  all  the  water  is  driven  off,  and  the  resulting  anhydrous 
salt  is  blue.  The  pink  aqueous  solution  also  turns  blue  on  heating, 
but  regains  its  original  colour  on  cooling.  Advantage  is  taken  of 
this  behaviour  in  preparing  the  so-called  sympathetic  ink.  Writing 
done  with  an  aqueous  solution  of  cobalt  chloride  is  practically 
invisible,  but  when  the  paper  is  warmed  the  blue  anhydrous  salt  is 
formed,  and  the  writing  becomes  visible.  It  fades  away  again  on 
cooling,  owing  to  rehydration  of  the  salt. 

Cobaltous  Sulphate,  CoSO4,7H2O,  prepared  by  the  usual 
methods,  occurs  in  dark  red  monoclinic  crystals,  isomorphous  with 
FeSO4,7H2O.  It  can  be  dehydrated  by  heating,  and  the  anhydrous 
salt  is  also  red.  It  forms  double  salts  with  the  alkali  sulphates  ; 
for  example,  CoSO4,K2SO4,6H2O,  isomorphous  with  the  correspond- 
ing iron  salt. 

Cobaltous  Nitrate,  Co(NO3)3,6H2O,  forms  red  hygroscopic 
prisms,  and  is  very  soluble  in  water.  It  gives  cobaltic  oxide  when 
cautiously  heated. 

Cobaltous  Sulphide,  CoS,  is  obtained  as  a  black  amorphous 
precipitate  by  the  addition  of  ammonium  sulphide  to  a  solution  of  a 
cobalt  salt.  When  hydrogen  sulphide  is  passed  through  a  cobalt 
solution  acidified  with  hydrochloric  acid  no  precipitate  is  formed, 
but  the  sulphide  obtained  by  precipitation  in  alkaline  solution  is 
practically  insoluble  in  dilute  hydrochloric  acid.  This  curious  be- 
haviour is  not  entirely  understood. 

Sulphides  of  the  formulae  Co2S3,  Co3S4  and  CoS2  have  also  been 
described. 

COBALTIC  SALTS 

As  already  mentioned,  simple  salts  containing  trivalent  cobalt  are 
highly  unstable. 

Cobaltic  Sulphate,  Co2(SO4)3,  is  formed  in  solution  at  the 
anode  when  a  concentrated  solution  of  cobaltous  sulphate,  CoSO4, 
is  electrolyzed.  It  has  been  obtained  in  the  solid  form  as 
Co2(SO4)3,i8H2O.  The  aqueous  solution  is  dark  green. 

Many  complex  compounds  containing  trivalent  cobalt  are  known, 
and  will  now  be  briefly  referred  to. 


COBALT  567 

i 
COMPLEX  SALTS  CONTAINING  COBALT 

Complex  Cyanides — The  complex  salts,  K4Co(CN)6  and 
K3Co(CN)6,  corresponding  with  potassium  ferro-  and  ferricyanide, 
are  known. 

When  excess  of  potassium  cyanide  is  added  to  a  cobaltous  salt  in 
the  cold,  a  red  solution  is  obtained,  which,  on  evaporation  at  room 
temperature,  yields  potassium  cobaltocyanide  in  violet  crystals  : 

CoCl2  +  6KCN->K4Co(CN)6  +  2KCl. 

When  the  red  solution  is  boiled  it  becomes  decolorized,  hydrogen  is 
given  off,  and  on  evaporation  the  stable  potassium  cobalticyanide, 
K3Co(CN)6,  is  obtained  in  colourless  crystals. 

Complex  Nitrite — When  potassium  nitrite  and  acetic  acid  are 
added  to  the  solution  of  a  cobaltous  salt,  the  red  colour  disappears, 
and  in  course  of  time  the  cobalt  is  completely  precipitated  in  the 
form  of  a  yellow  crystalline  substance,  of  the  formula  K3Co(NO2)6, 
potassium  cobaltinitrite.  The  complex  anion  has  the  formula 
Co(NO2)G"'.  The  salt  can  be  classified  from  the  point  of  view  of 
Werner's  theory  (p.  577). 

Cobaltammines — When  to  a  solution  ot  a  cobaltous  salt  excess 
of  ammonia  is  added,  and  the  resulting  solution  is  treated  with  an 
oxidizing  agent,  e.g.  the  oxygen  of  the  air,  a  great  variety  of  complex 
salts  are  obtained,  the  empirical  composition  of  which  corresponds 
with  that  of  a  cobaltic  salt  associated  with  3,  4,  5  or  6  NH3  molecules. 
The  compounds,  however,  unlike  the  cobaltic  salts,  are  quite  stable, 
and  therefore  do  not  contain  simple  Co'"  ions  ;  further,  in  some  of 
them  only  part  of  the  anion  shows  the  reactions  characteristic  of  it, 
for  example,  only  part  of  the  chlorine  is  precipitated  by  silver  nitrate. 
In  these  and  other  respects  the  compounds  resemble  the  chromium 
ammonia  compounds,  and  they  can  be  represented  as  follows  : — 

[Co(NH3)6]-Cl3  ;  [Co(NH3)6Cl]"Cl2;  [Co(NH3)4Cl2]'Cl ; 

[Co(NH3)3Cl3]. 

The  complex  enclosed  in  the  square  brackets  is  the  cation.  In  the 
first  compound  it  is  trivalent,  and  all  the  chlorine  atoms  are  ionised  ; 
in  the  second  compound  the  cation  is  divalent,  and  in  the  third 
univalent.  Finally,  in  the  fourth  compound  it  consists  of  a  non- 
ionised  complex,  which  gives  none  of  the  ordinary  reactions  for  cobalt 
and  chlorine  ions.  The  behaviour  of  these  compounds  is  satisfac- 
torily accounted  for  on  Werner's  theory  (p.  577). 


568     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Tests  for  Cobalt — All  cobalt  salts  communicate  an  intense 
blue  colour  to  a  borax  bead.  The  formation  of  a  black  precipitate 
with  ammonium  sulphide,  insoluble  in  dilute  hydrochloric  acid, 
coupled  with  the  fact  that  hydrogen  sulphide  in  acid  solution  causes 
no  precipitate,  is  characteristic  of  nickel  and  cobalt  salts.  The 
reactions  whereby  nickel  and  cobalt  salts  may  be  distinguished  are 
referred  to  under  nickel. 

NICKEL 

Symbol,  Ni.     Atomic  weight,  58.68. 

Occurrence — Nickel  is  found  free  in  meteorites.  In  combina- 
tion it  is  met  with  in  kupfernickel  or  niccolite,  NiAs,  in  nickel  glance^ 
NiSAs,  as  a  silicate  in  garnierite,  H.2(Mg,Ni)SiO4,2H2O  (New 
Caledonia),  and  as  sulphide,  along  with  the  sulphides  of  iron  and 
copper,  in  Canada.  It  is  now  obtained  almost  exclusively  from  the 
two  latter  sources. 

Preparation  of  Metal — The  metal  may  be  obtained  from  the 
oxide  by  reducing  with  carbon  or  hydrogen,  or  by  the  Goldschmidt 
process  with  aluminium  powder.  Very  pure  nickel  is  now  obtained 
commercially  by  Mond's  process,  which  depends  on  the  intermediate 
formation  of  nickel  carbonyl,  Ni(CO)4  (p.  335).  Nickel  oxide  (obtained 
by  roasting  nickel  ores)  is  heated  and  carbon  monoxide  passed  over 
it,  whereby  reduction  to  the  metal  first  takes  place,  the  latter  then 
uniting  with  carbon  monoxide  to  form  the  volatile  nickel  carbonyl. 
The  vapour  of  the  carbonyl  is  then  passed  through  tubes  at  a  higher 
temperature  (about  200°  C.),  whereby  it  is  decomposed,  and  nickel  is 
deposited  in  coherent  form. 

Properties — Nickel  is  a  silvery-white,  lustrous  metal  of  density 
8.8  to  9.  i  ;  it  melts  at  1435°.  **  *s  malleable  and  ductile,  and  also 
hard  and  tenacious.  It  is  magnetic,  but  loses  this  property  on 
heating  to  250°.  Like  iron,  it  becomes  passive  on  treatment  with 
oxidizing  agents. 

Nickel  is  practically  unaffected  by  exposure  to  air.  It  is  attacked 
very  slowly  by  hydrochloric  or  sulphuric  acid,  but  dissolves  fairly 
readily  in  dilute  nitric  acid,  nickel  nitrate,  Ni(NO3)2,  being  formed. 

Alloys  Of  Nickel — Nickel  is  a  constituent  of  many  useful 
alloys.  German  silver  usually  contains  50  per  cent,  of  copper,  25  per 
cent,  of  nickel,  and  25  per  cent,  of  zinc,  but  the  proportions  vary  to 
some  extent.  The  nickel  coins  in  use  on  the  Continent  and  in  the 
United  States  are  made  of  an  alloy  of  75  per  cent,  of  copper  and 
25  per  cent,  of  nickel.  It  is  noteworthy  that  in  spite  of  the  high 


.    NICKEL  569 

proportion  of  copper  this  alloy  is  white.  Nickel  steel  has  already 
been  mentioned  (p.  558).  Further,  on  account  of  its  permanence  in 
the  air,  nickel  is  used  as  a  coating  for  other  metals.  Nickel  plating, 
as  it  is  called,  is  done  electrolytically,  a  solution  of  nickel  sulphate  or 
nickel  ammonium  sulphate  being  used  as  electrolyte  and  a  nickel 
plate  as  anode. 


OXIDES  AND  HYDROXIDES  OF  NICKEL 

Three  oxides  of  nickel  are  known:  nickelous  oxide,  NiO,  nickelic 
oxide,  Ni2O3,  and  nickelo-nickelic  oxide,  Ni3O4.  Only  one  series  of 
nickel  salts  is  known,  corresponding  with  nickelous  oxide,  NiO. 
Nickelic  oxide  behaves  towards  acids  as  a  peroxide. 

Nickelous  Oxide,  NiO,  is  obtained  as  a  grayish-green  powder 
by  heating  nickelous  hydroxide  out  of  contact  with  air. 

Nickelous  Hydroxide,  Ni(OH)2,  is  obtained  as  an  apple- 
green  precipitate  when  an  alkali  hydroxide  is  added  to  the  solution  of 
a  nickel  salt.  The  hydroxide  is  insoluble  in  excess  of  alkali  hydroxide, 
but  dissolves  in  ammonia  to  form  a  blue  solution.  Unlike  the  corre- 
sponding cobalt  solution  (p.  565),  the  ammoniacal  solution  of  nickel 
hydroxide  does  not  absorb  oxygen  from  the  air.  It  contains  complex 
ions  of  the  type  Ni(NH4)a;,  in  which  the  nickel  is  divalent.  They  are 
much  less  stable  than  the  cobaltammines,  which,  however,  are  derived 
from  tervalent  cobalt. 

Nickelic  Oxide,  Ni2O3,  is  obtained  as  a  black  powder  by 
decomposing  nickelous  nitrate  by  heat  at  as  low  a  temperature  as 
possible.  When  heated  with  hydrochloric  acid  it  forms  nickelous 
chloride  and  chlorine,  whilst  with  sulphuric  acid  it  gives  nickelous 
sulphate  and  oxygen  : 


Ni2O3+2H2SO4->2NiSO4  +  2 

It  therefore  behaves  as  a  peroxide,  and  appears  to  have  no  basic 
^properties. 

Nickelic  Hydroxide,  Ni(OH)3,  is  obtained  as  a  black  pre- 
cipitate by  adding  an  alkali  hypochlorite  or  a  solution  of  bleaching 
powder  to  the  solution  of  a  nickel  salt.  Chemically  it  behaves  like 
the  corresponding  oxide. 

Nickelo-nickelic  Oxide,  Ni3O4,  is  said  to  be  formed  as  a 
gray  mass  when  moist  oxygen  is  passed  over  nickel  chloride 


570     A   TEXT-BOOK   OF   INORGANIC    CHEMISTRY 

heated  at  400°.     On  heating  strongly  oxygen  is  given  off  and  the 
monoxide,  NiO,  is  formed. 

SALTS  OF  NICKEL 

As  already  mentioned,  the  salts  of  nickel  are  derived  exclusively 
from  the  monoxide,  NiO,  and  therefore  contain  the  bivalent  Ni"  ion, 
which  is  green.  The  hydrated  salts  and  solutions  are  green,  the 
anhydrous  salts  yellow  or  brown. 

Nickel  Chloride,  NiCl2,  is  obtained  by 'dissolving  the  oxide  in 
hydrochloric  acid,  and  separates  from  solution  with  6H2O  in  green 
monoclinic  prisms.  When  heated  the  anhydrous  salt,  which  is 
yellow,  is  obtained.  The  latter  absorbs  gaseous  ammonia,  forming 
the  compound  Ni(NH3)6Cl2,  which  separates  from  aqueous  solution 
in  blue,  octahedral  crystals. 

Nickel  Sulphate,  NiSO4,  obtained  by  dissolving  the  metal 
or  the  oxide  in  dilute  sulphuric  acid,  separates  from  solution  below 
20°  with  yH2O  in  green,  rhombic  prisms,  isomorphous  with  the 
corresponding  ferrous  and  other  sulphates.  At  a  higher  temperature 
(30  to  40°)  green,  tetragonal  crystals  of  the  formula  NiSO4,6H2O  are 
obtained.  When  heated  at  100°  it  loses  6H2O,  above  300°  it  becomes 
anhydrous.  With  excess  of  ammonia  a  violet  solution  is  obtained,  from 
which  on  evaporation  violet  crystals  of  the  formula  NiSO4,4NH3,2H2O 
are  obtained.  This  salt  is  presumably  [Ni(NH3)4(H2O)2]SO4  ;  many 
other  instances  in  which  water  displaces  ammonia  in  such  complexes 
are  known. 

With  alkali  sulphates,  nickel  sulphate  forms  double  salts,  such  as 
NiSO4,(NH4)2SO4,6H2O,  isomorphous  with  the  corresponding  iron 
salt.  This  salt  is  largely  used  in  electro-plating. 

Nickel  Sulphides — Nickelous  sulphide,  NiS,  is  obtained  as  a 
black  precipitate  by  adding  ammonium  sulphide  to  a  solution  of  a 
nickel  salt.  Like  the  corresponding  cobalt  salt,  it  is  practically 
insoluble  in  dilute  hydrochloric  acid,  buj:  is  soluble  to  some  extent  in 
ammonium  sulphide  solution. 

The  sulphides  Ni3S4  and  NiS2  have  also  been  described. 

Nickel  Cyanide,  Ni(CN)2,  is  obtained  as  a  green  precipitate  by; 
adding  potassium  cyanide  to  a  solution  of  a  nickel  salt.  The  pre- 
cipitate dissolves  in  excess  of  cyanide  to  form  the  compound  K2Ni(CN)4, 
which  can  be  obtained  as  the  monohydrate,  in  yellow  crystals,  by 
evaporating  the  solution.  Unlike  the  nearly  analogous  compound 
containing  divalent  cobalt,  it  is  stable  on  boiling,  but  is  much  less 
stable  than  potassium  cobalticyanide  (see  tests  for  nickel). 


THE   IRON   SUB-GROUP—GROUP   VIII 


Tests  for  Nickel — Nickel  salts  colour  the  borax  bead  yellowish 
brown  in  the  oxidizing  flame.  The  behaviour  of  the  hydroxide  and 
of  the  sulphide  is  made  use  of  in  detecting  nickel.  Nickel  and  corral  t 
compounds  are  distinguished  by  the  facts  (i)  that  only  the  latter  give 
a  yellow  precipitate  with  acetic  acid  and  potassium  nitrite ;  (2)  when 
sodium  hypochlorite  or  hypobromite  is  added  to  a  solution  of  a  nickel 
salt  in  excess  of  potassium  cyanide  a  black  precipitate  of  nickelic 
hydroxide  is  obtained,  whereas  a  solution  of  a  cobalt  salt  in  potas- 
sium cyanide  after  boiling  gives  no  precipitate  under  these  conditions. 

Summary  of  Metals  of  Iron  Sub-group — Corresponding 
with  the  fact  that  the  members  of  this  family  differ  only  slightly  in 
atomic  weight,  their  physical  constants  do  not  differ  very  greatly. 
This  is  shown  in  the  accompanying  table : — 


Iron. 

Cobalt. 

Nickel. 

Atomic  weight 

55-85 

58.97 

58.68 

Density    .... 

7.86 

8.5 

8.9 

Melting-point  .         . 

1500° 

1460° 

1435° 

Atomic  volume 

7.1 

6-9 

6.6 

Further,  all  three  metals  are  grayish-white,  magnetic,  and  readily 
become  passive  when  treated  with  concentrated  nitric  acid. 

From  the  chemical  point  of  view,  all  three  metals  form  oxides  of  the 
type  MO,  which  are  strongly  basic  and  give  rise  to  stable  series  of 
salts.  Further,  all  have  oxides  of  type  M2O3,  but  corresponding  salts 
are  only  known  for  iron  and  cobalt ;  nickel  salts  in  which  the  metal 
is  trivalent  are  unknown.  Only  iron  forms  salts  corresponding  with 
a  (hypothetical)  acidic  oxide,  FeO3,  so  that  cobalt  and  nickel  are  more 
metallic  than  iron.  The  latter  two  elements  resemble  each  other  very 
closely  in  their  compounds  ;  the  chief  difference  is  the  greater  ten- 
dency of  cobalt  to  act  as  a  trivalent  element,  as  shown,  for  instance,  in 
the  cobaltammines.  Iron  has  comparatively  little  tendency  to  enter 
into  the  formation  of  complex  ions,  and  in  this  important  respect  differs 
from  cobalt  and  nickel. 

As  regards  comparison  with  members  of  other  groups,  there  is 
a  close  analogy,  especially  in  the  case  of  iron,  with  chromium,  man- 
ganese and  copper.  Reference  to  this  has  already  been  made  in 
connexion  with  the  latter  metals. 


CHAPTER    XXXVI 
THE   PLATINUM   SUB-GROUP  (GROUP   VIII) 

Ruthenium,  Ru=ioi.7         Rhodium,  Rh  =  ro2.6        Palladium,  Pd  = 
Osmium,  Os       —190.9         Iridium,  Ir     —193.1         Platinum,  Pt    =195.2 

THE  remaining  elements  of  the  eighth  group  are  naturally  divided 
into  two  smaller  groups  containing  three  elements  each.  Just 
as  iron,  cobalt  and  nickel  differ  only  slightly  in  atomic  weight  and 
in  density  (p.  571),  so  the  atomic  weights  of  the  triad  ruthenium, 
rhodium  and  palladium  vary  from  101.7  to  106.7,  and  their  densities 
from  11.4  to  12.26;  the  atomic  weights  of  the  triad  osmium,  iridium 
and  platinum  vary  from  190.9  to  195.2,  and  their  densities  from  21.5 
to  22.5.  As  in  the  remainder  of  the  periodic  table,  however,  there  is 
also  considerable  resemblance  between  the  elements  in  the  same 
vertical  series  ;  thus  the  pairs  ruthenium  and  osmium,  rhodium  and 
iridium,  and  palladium  and  platinum,  show  considerable  chemical 
analogy. 

The  six  metals  occur  in  the  free  state  in  nature,  associated  in 
alluvial  deposits  in  the  so-called  platinum  ore,  which  contains  70  to 
80  per  cent,  of  platinum,  5  to  8  per  cent,  of  iridium,  and  a  smaller 
proportion  of  the  other  metals.  Gold  and  copper  are  also  generally 
present.  The  platinum  ores  are  found  chiefly  in  the  Urals  and  in 
Brazil.  They  are  separated  from  the  sand,  etc.,  by  washing,  advan- 
tage being  taken  of  their  relatively  great  density.  As  platinum  is  the 
only  important  member  of  the  group,  the  others  will  be  described 
very  briefly. 

Ruthenium  (Ru— 101.7)  *s  a  steel-gray  metal  of  density  12.26;  it  melts  about 
2000°.  When  heated  in  the  air  it  becomes  covered  with  a  brown  film  of  the 
dioxide,  and  when  heated  in  a  current  of  oxygen  the  dioxide,  RuO2,  is  formed. 
It  is  only  slightly  attacked  by  aqua  regia,  and  is  chiefly  obtained  from  the  alloy 
osmiridium,  which  remains  when  platinum  ores  are  treated  with  aqua  regia. 

Ruthenium  forms  the  oxides,  Ru2O3,  RuO2  and  RuO4,  and  also  salts  corre- 
sponding with  the  acidic  oxides,  RuO3  and  Ru2O7.  A  salt  corresponding  with 
the  trioxide  is  potassium  ruthenate,  K2RuO4 ;  when  diluted  it  decomposes  into 
potassium  perruthenate,  KRuO4,  and  an  oxide  of  ruthenium  (cf.  manganates  and 

S7» 


THE   PLATINUM   SUB-GROUP—GROUP   VIII     573 

permanganates).  The  most  stable  ruthenium  salts  contain  tervalent  ruthenium, 
but  compounds  in  which  the  metal  is  bivalent  and  quadrivalent  are  also  known. 

Osmium  (03=190.9)  is  a  bluish-white  crystalline  metal ;  its  density  is  22.5,  and 
it  melts  about  2300°.  On  account  of  its  difficult  fusibility,  it  is  used  as  a  filament 
in  incandescent  lamps.  On  heating  in  air,  it  burns  readily  to  the  easily  volatile 
tetroxide,  which  is  very  poisonous,  and  has  a  highly  injurious  effect  on  the  eyes. 
The  metal  is  scarcely  attacked  by  aqua  regia. 

Osmium  forms  four  oxides,  OsO,  08203,  OsO2andOsO4  ;  and  salts,  theosnpates, 
are  known,  derived  from  the  unknown  acidic  oxide,  OsO3.  The  best  known  is 
potassium  osmate,  K2OsO4,2H2O,  which  occurs  in  garnet  red  octahedral  crystals. 
The  most  remarkable  compound  of  osmium  is  the  tetroxide,  OsO4,  usually  known 
as  osmic  acid,  although  it  does  not  appear  to  possess  acidic  properties.  It  occurs 
in  colourless,  lustrous  needles,  which  melt  about  40° ;  the  liquid  boils  about  100°. 
It  is  used  for  staining  specimens  in  biological  work,  the  oxide  being  reduced, 
most  readily  by  fats,  to  metallic  osmium. 

It  is  interesting  to  note  that,  corresponding  with  their  position  in  the  eighth 
group  of  the  periodic  table,  ruthenium  and  osmium  have  a  maximum  valency 
of  8. 

Rhodium  (Rh=  102.6)  is  a  silver-white,  lustrous,  malleable  metal  of  density  12.1, 
which  fuses  at  a  higher  temperature  than  platinum.  It  is  not  attacked  by  aqua 
regia. 

Three  oxides  of  rhodium  are  known,  RhO,  Rh2O3  and  RhO2.  The  most  stable 
salts  are  derived  from  the  oxide  Rh2O3.  Unlike  ruthenium  and  osmium,  rhodium 
forms  no  compounds  derived  from  an  acidic  oxide.  The  rhodium  salts  usually 
form  red  aqueous  solutions,  hence  the  name.  The  trichloride,  RhCl3,  forms 
stable  compounds  with  the  alkali  chlorides,  e.g.  Na3RhCl6  (or  RhCl3,3NaCl). 

Iridium  (Ir=i93.i)  is  a  white,  lustrous,  brittle  metal  of  density  22.4.  Next  to 
osmium,  it  has  the  highest  melting-point  of  the  platinum  metals.  In  the  free 
condition  it  is  not  attacked  by  aqua  regia,  but  when  alloyed  with  much  platinum 
it  is  dissolved  to  some  extent.  On  account  of  their  high  melting-point  and  resist- 
ance to  many  reagents,  alloys  of  platinum  and  iridium  are  used  for  various 
purposes. 

Iridium  forms  two  basic  oxides,  Ir2O3  and  IrO2,  from  which  two  series  of  salts 
are  derived.  The  chlorides,  IrCl3  and  IrCl4,  form  complex  compounds  with  the 
alkali  chlorides  of  the  types  K3IrCl6  or  IrCl3,3KCl,  and  K2IrCl6  or  IrCl4,2KCl, 
which  correspond  with  the  better  known  platinum  compounds. 

Palladium  (Pd= 106.7)  is  found  in  platinum  ore  and  also,  alloyed  with  gold,  in 
certain  parts  of  South  America.  It  is  a  silvery-white,  lustrous,  malleable  metal  of 
density  11.9,  and  melts  at  1546°.  The  most  remarkable  property  6f  the  metal  is  its 
power  of  absorbing  hydrogen,  already  referred  to  (p.  39).  The  freshly  ignited 
metal  absorbs  600  times  its  own  volume  of  hydrogen  under  ordinary  conditions, 
and  when  made  the  cathode  of  an  electrolytic  cell  at  which  hydrogen  is  being 
generated  it  takes  up  about  loco  volumes  of  the  gas.  The  hydrogen  is  completely 
expelled  on  heating  to  redness.  It  does  not  appear  to  be  chemically  combined 
with  the  palladium,  but  is  present  in  solid  solution.  The  hydrogen  absorbed  in 
palladium  has  powerful  reducing  properties,  due  partly  to  the  catalytic  effect  of 
the  palladium  on  the  speed  of  reduction  and  partly  to  the  very  condensed  condi- 
tion of  the  hydrogen. 

Two  oxides  of  palladium,  PdO  and  PdO2,  are  known.    Well-defined  palladium 


574     A   TEXT-BOOK    OF   INORGANIC   CHEMISTRY 

salts  are  derived  exclusively  from  the  first-mentioned  oxide.  When  palladium 
is  dissolved  in  aqua  regia,  palladium  tetrachloride  is  present  in  solution  as 
H2PdCl6  or  PdCl4,2HCl;  it  has  not  been  obtained  in  the  free  condition.  When 
the  solution  is  boiled,  chlorine  is  given  off  and  palladious  chloride,  PdCLj  (or 
rather  H2PdCl4),  is  formed. 

PLATINUM 

Symbol,  Pt.     Atomic  weight,  195.2. 

Preparation  of  Metal — The  platinum  ore  is  treated  with 
dilute  aqua  regia,  whereby  platinum,  palladium,  rhodium  and  indium 
are  dissolved  as  the  higher  chlorides.  The  solution  is  evaporated  to 
dryness  and  the  fused  mass  heated  at  125°,  at  which  temperature  the 
palladium  and  rhodium  salts  are  reduced  to  lower  chlorides.  The 
residue  is  extracted  with  water,  acidified  with  hydrochloric  acid 
and  ammonium  chloride  added,  when  ammonium  platinic  chloride, 
PtCl4,2NH4Cl,  is  precipitated  in  yellow  crystals  (see  below).  From 
the  solution  the  corresponding  iridium  salt,  IrCl4,2NH4Cl,  which  is 
more  soluble,  is  obtained  by  concentrating  the  solution.  The  double 
platinum  salt,  on  ignition,  yields  the  metal  in  spongy  form.  Platinum 
is  obtained  in  lustrous,  coherent  form  by  fusing  the  spongy  metal  in  a 
lime  crucible  by  means  of  the  oxyhydrogen  flame. 

Properties — Platinum  is  a  silvery-white,  malleable  and  ductile 
metal  of  density  21.42.  It  fuses  at  1750°,  and  therefore  does  not  melt 
in  the  Bunsen  flame,  but  readily  fuses  in  the  oxyhydrogen  flame  ;  it 
can  be  welded  at  a  white  heat.  Massive  platinum  does  not  oxidize 
in  the  air  even  at  high  temperatures.  The  metal  is  insoluble  in  any 
single  a^id,  but  is  readily  dissolved  by  aqua  regia  with  formation  of 
platinum  tetrachloride.  It  forms  alloys  with  lead  and  antimony, 
and  certain  other  heavy  metals,  and  combines  with  carbon,  silicon 
and  phosphorus  to  form  brittle  alloys.  Further,  it  is  attacked  by 
fused  alkalis,  cyanides,  nitrates  and  sulphides.  These  facts  should, 
be  borne  in  mind  in  using  platinum  utensils  (crucibles,  wire,  etc.), 
which,  on  account  of  the  slight  chemical  activity  of  the  metal,  are 
largely  employed  in  the  laboratory. 

Platinum  condenses  gases,  such  as  oxygen  and  hydrogen,  on  its 
surface,  especially  in  the  finely  divided  (spongy)  form.  Partly  on 
this  account,  the  metal  has  a  remarkable  effect  in  accelerating  many 
chemical  reactions,  such  as  the  combination  of  inflammable  gases 
(hydrogen,  marsh  gas,  benzene,  etc.),  with  oxygen.  A  mixture  of 
oxygen  and  hydrogen  can  be  brought  to  explosion  at  room  tempera- 
ture by  bringing  it  in  contact  with  finely  divided  platinum.  In  recent 


PLATINUM 


575 


years  it  has  been  found  that  the  finely  divided  metal  oxidizes  much 
more  readily  than  was  formerly  supposed,  and  it  is  not  improbable 
that  the  intermediate  formation  of  an  active  platinum  oxide  plays  a 
part  in  these  phenomena. 

*  Spongy  platinum,,  as  already  explained,  is  obtained  by  heating 
ammonium  platinic  chloride  ;  platinum  black,  another  finely  divided 
form  of  the  metal,  is  obtained  by  electrolysis  of  a  solution  of  platinic 
chloride,  or  by  adding  finely  divided  zinc  to  the  latter  solution. 
These  forms  of  platinum  have  very  high  absorptive  power  for 
hydrogen;  platinum  black  absorbs  about  100  times  its  volume  of 
the  gas.  Platinum  black  also  absorbs  about  100  times  its  volume 
of  oxygen,  but  in  this  case  partial  chemical  combination  probably 
occurs. 

Platinum  Compounds — Two  series  of  platinum  compounds 
are  known  ;  platinous  compounds,  e.g.  PtCl2,  in  which  the  metal  is 
divalent,  and  platinic  compounds,  e.g.  PtCl4,  in  which  the  platinum 
is  quadrivalent. 

Oxides  and  Hydroxides  of  Platinum— Platinous  oxide, 
PtO,  is  obtained  as  a  gray  powder,  and  platinic  oxide,  PtO2,  as  a 
black  powder,  by  cautiously  heating  the  corresponding  hydroxides. 
When  strongly  heated,  they  lose  oxygen  and  yield  the  metal. 

Platinous  Hydroxide,  Pt(OH)2,  is  obtained  as  a  dark  pre- 
cipitate by  the  action  of  sodium  hydroxide  solution  on  platinous 
chloride.  It  is  a  typical  base  and  dissolves  in  acids  to  form  platinous 
salts. 

Platinic  Hydroxide,  Pt(OH)4,  is  obtained  as  the  dihydrate, 
Pt(OH)4,2H2O,  or  H2Pt(OH)6  (platinic  acid),  by  boiling  platinic 
chloride  with  sodium  hydroxide  solution,  cobling  and  neutralizing  with 
acetic  acid.  It  is  a  yellow  powder,  which  on  heating  turns  brown  and 
then  black.  It  has  weak  basic  and  also  weak  acidic  properties  ;  with 
potassium  hydroxide  it  gives  potassium  platinate,  K2PtO3. 

Platinous  Chloride,  PtCl2,  is  obtained  by  heating  platinic 
chloride,  PtCl4,  to  300°.  It  forms  a  grayish-green  powder,  insoluble 
in  water  but  soluble  in  hydrochloric  acid.  With  the  alkali  chlorides 
it  forms  double  salts,  e.g.  PtCl2,2KCl  and  PtCl2,2NH4Cl,  both  of 
which  are  red,  well  crystallized  salts.  On  account  of  their  behaviour, 
it  is  assumed  that  these  salts  contain  the  complex  anion  PtCl4", 
and  they  are  therefore  termed  chloroplatinites,  (NH4)2PtCl4,  and 
K2PtCl4,  derived  from  chloroplatinous  acid,v  H2PtCl4.  The  latter 
compound  is  presumably  present  in  the  solution  of  the  dichloride  in 
hydrochloric  acid. 


576    A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 

Platinum  Trichloride,  PtCl3,  is  obtained  as  a  green  powder 
by  prolonged  heating  of  platinic  chloride  in  dry  chlorine  at  320° 
(Wohler  and  Martin,  1909). 

Platinic  Chloride,  PtCl4,  is  obtained  by  the  direct  action  of 
chlorine  on  platinum  at  a  high  temperature.  It  is  obtained  more 
conveniently,  however,  by  dissolving  platinum  in  aqua  regia,  evaporat- 
ing to  dryness,  and  heading  the  residue  in  a  current  of  hydrogen 
chloride.  It  separates  from  aqueous  solution  in  large  red  crystals 
as  PtCl4,5H2O.  When  platinum  is  dissolved  in  aqua  regia  and 
the  solution  is  evaporated  to  dryness  several  times  with  hydro- 
chloric acid  till  all  the  nitric  acid  is  expelled,  the  compound 
PtCl4,2HCl  or  H2PtCl6,  usually  termed  chloroplatinic  acid,  .  is 
obtained.  It  separates  from  aqueous  solution  in  brownish-red 
prisms  with  6H2O. 

Platinic  chloride  forms  double  compounds  with  the  alkali  chlorides, 
which  may  be  regarded  as  being  derived  from  chloroplatinic  acid  by 
displacing  the  hydrogen  by  metals.  Owing  to  their  different  solu- 
bilities in  water,  they  are  used  in  analytical  chemistry. 

Potassium  Platinichloride,  K2PtCl6,  obtained  as  a  yellowish 
crystalline  precipitate  by  adding  a  potassium  salt  to  a  solution  of 
chloroplatinic  acid,  is  only  slightly  soluble  in  water.  At  o°  100  grains 
of  water  dissolve  0.70  grams,  at  20°  1.12  grams,  and  at  40°  1.76 
grams  of  the  salt.  Potassium  platinichloride  is  almost  insoluble 
in  alcohol.  The  corresponding  ammonium  salt  is  slightly,  the 
sodium  salt  readily  soluble  in  water,  and  on  this  is  based  a 
method  of  separating  potassium  and  sodium  in  analysis. 

Platinocyanides  and  Platinicyanides — Corresponding 
with  the  platinochlorides,  platinocyanides,  derived  from  platino- 
cyanic  acid,  H2Pt(CN)4,  are  known.  The  acid  and  the  salts  are 
characterized  by  remarkable  differences  in  colour  depending  on  the 
proportion  of  water  with  which  they  are  associated,  and  further, 
the  crystals  show  different  colours  in  different  directions.  Platini- 
cyanides, derived  from  platinicyanic  acid,  H2Pt(CN)6,  are  also 
known. 

Platinum  Sulphides — Platinous  sulphide,  PtS,  and  platinic 
sulphide,  PtS2,  are  obtained  as  black  precipitates  by  passing 
hydrogen  sulphide  through  solutions  of  the  respective  chlorides. 
They  are  insoluble  in  hydrochloric  acid,  but  dissolve  in  solutions 
of  alkali  sulphides,  forming  the  salts. 

Platinum  Ammonia  Compounds— Like  salts  of  cobalt,  chromium  and  other 
metals,  platinum  salts  form  compounds  with  ammonia  which  can  be  classified 


MODERN  VIEWS  ON  VALENCY  577 

on  the  basis  of  Werner's  theory  of  valency  (see  below).     Thus  the  following 
series  of  compounds  containing  divalent  platinum  is  known  : 


[  Pt(NH3)l  ]"0,  ;     [Pt  ("£*]•  C,  ;     [Pt  (NHj),  ]  . 


All  the  compounds  contain  a  special  group,  PtA4,  the  character  of  which  varies 
with  the  nature  of  the  groups  A.  In  the  first  two  members  it  is  basic,  in  the 
middle  member  it  is  neutral,  and  in  the  two  last  members  it  is  acidic.  The 
further  behaviour  of  these  compounds  will  be  understood  on  the  basis  of  the  ex- 
planations given  in  the  following  sections. 

The  compound  Pt'   Q       exists  in  two  entirely  distinct  forms  which  differ  in 

many  of  their  reactions.  An  explanation  of  this  case  of  isomerism  on  the  basis  of 
a  different  arrangement  of  the  atoms  in  space  is  given  below  (p.  584). 

Another  series  of  compounds,  containing  quadrivalent  platinum,  is  known  : 
[Pt(NH8)6]--Cl4;  [Pt(NH3)5Cl]"'Cl8;   [Pt(NH3)4Cla]"  C12  ;    [Pt(NH8)Cl8]-  Cl  ; 

[Pt(NH3)2Cl4]  ;  [Pt(NH3)Cl6]'  K  ;  [PtClJ"  Ka. 

All  these  compounds  contain  a  special  group,  PtA6,  which  in  the  first  members 
is  electro-positive,  but  gradually  diminishes  in  valency  as  the  number  of  Cl  atoms 
in  it  increases,  till  it  becomes  electrically  neutral  (compound  five),  and  finally 
electro-negative.  Werner's  theory  of  valency,  which  affords  a  satisfactory  repre- 
sentation of  the  behaviour  of  such  compounds,  will  be  considered  in  a  later  section 
(P-  58o). 

Valency  —  Valency  has  been  defined  (p.  130)  as  the  number  of 
atoms  of  hydrogen  or  of  a  halogen  with  which  an  atom  of  the  element 
in  question  can  combine  to  form  a  molecule.  An  alternative  de- 
finition is  obtained  from  the  expression 

Atomic  Weight      _ 
Chemical  Equivalent  = 

already  fully  explained  (p.  131).  If  the  chemical  equivalent  is  deter- 
mined by  reference  to  the  same  compound  as  has  been  used  to  establish 
the  valency  according  to  the  first  definition,  the  valency  deduced 
according  to  the  two  methods  is  the  same.  If  different  compounds 
are  used,  however,  this  is  not  necessarily  the  case. 

The  reason  why  univalent  elements,  such  as  hydrogen  and  the 
halogens,  are  used  to  determine  valency  in  preference  to  polyvalent 
elements  is  because  the  atoms  of  the  latter  elements  may  be  linked 
together  by  one  or  more  valencies.  Thus  for  sulphur  trioxide,  SO3, 
the  following  graphic  formulae  are  possible  :  — 

/&  ,0        /Ox 

O  —  <*//      .  n  —  c/   i  .  c./      \n 
—  O^  ,     U  —  Ov  ,     O.  ,\J 

^o          No     No/ 

in  which  sulphur  is  bivalent,  quadrivalent,  and  sexavalent  respectively 
37 


57«     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

and  it  is  not  possible  without  further  evidence  to  decide  between 
them.  On  the  other  hand,  the  existence  of  sulphur  hexafluoride,  SF6 
(p.  297)  is  fairly  conclusive  evidence  that  the  maximum  valency  of 
sulphur  is  at  least  six. 

The  valency  of  an  element  may  also  be  deduced  from  the  number 
of  univalent  groups  with  which  an  atom  of  it  can  combine.  The 
compounds  formed  with  the  organic  groups  CH3  and  C2H5  (p.  346) 
are  particularly  useful  in  this  connection  (e.g.  Zn(CH3)2,  Sb(CH3)3, 
Pb(C2H5)4),  as  their  vapour  density  and  therefore  their  molecular 
weight  can  readily  be  determined. 

When  the  valency  of  an  element  has  been  established  by  the  method 
just  described,  the  information  thus  gained  is  used  to  obtain  informa- 
tion about  the  structure  of  more  complicated  compounds  by  writing 
their  graphic  formulae.  The  value  of  the  method  is,  however,  very 
much  lessened  by  the  fact  repeatedly  illustrated  in  the  course  of  the 
book,  that  most  elements  show  more  than  one  valency.  The  attempt 
was  made,  in  the  case  of  iron,  for  example,  to  save  the  hypothesis 
of  constant  valency  by  assuming  that  ferrous  and  ferric  chloride  were 
represented  by  the  respective  formulae 

Ck  /Cl  /Cl 

>Fe-Fe<         and  Fe— Cl 

Cl/  \C1  \C1 

but  this  view  became  untenable  when  it  was  shown  by  vapour  density 
determinations  that  at  high  temperatures  ferrous  chloride  possesses 
the  simple  formula  FeCl2. 

It  is  interesting  to  note  that  when  an  element  has  more  than  one 
valency,  these  valencies  often  differ  by  two.  Thus  sulphur  appears  to 
act  as  a  bivalent,  quadrivalent,  and  sexavalent  element,  and  the  chief 
valencies  of  nitrogen  and  phosphorus  are  three  and  five.  There  are, 
however,  many  exceptions  to  this  rule,  e.g.  ferrous  and  ferric  salts,  and 
the  existence  of  nitric  oxide,  which  undoubtedly  has  the  simple  formula 
NO  (p.  231). 

The  formulas  of  the  majority  of  chemical  compounds  can  be 
represented  graphically  on  the  assumption  that  the  valencies  of 
their  constituent  elements  are  those  derived  from  an  examination 
of  their  hydrogen  or  halogen  compounds.  Thus  the  chief  valencies 
of  phosphorus,  deduced  from  the  halogen  compounds,  are  three  and 
five,  and  in  accordance  with  this  the  graphic  formulas  of  the  principal 
oxides  are  usually  represented  as  follows  : — 

0>x  /& 

O  =  P  —  O  —  P-O,  phosphorous  oxide;  and        ^P-O  — P^ 


MODERN  VIEWS  ON  VALENCY  579 

phosphorous  pentoxide.  Similarly,  the  oxyacids  derived  from  these 
oxides  are  sometimes  assumed  to  contain  trivalent  and  quinquevalent 
phosphorus  respectively  (p.  255).  The  examination  of  a  number  of 
organic  phosphorous  compounds  has  shown,  however,  that  the  trivalent 
phosphorous  atom  has  a  tendency  to  become  oxidized  and  thus  become 

quinquevalent,  the  most   stable   arrangement  being     O  =  P  —  . 


formula  for  phosphorous  acid  may  therefore  be  written  O  =  P  —  OH,  and 

\OH 

we  have  already  seen  (p.  255)  that  this  formula  accords  best  with  its 
chemical  behaviour.  The  principal  valencies  of  lead,  deduced  from 
the  halogen  derivatives,  are  two  and  four,  and  in  accordance  with  this 
the  graphic  formula  of  lead  sesquioxide,  Pb2O3,  is  preferably  written 

•/°\ 

Pb(    .    /Pb  =  O,  and  not  as  O  =  Pb-O-Pb  =  O  in  which  the  lead 


atom  would  be  trivalent. 

It  must,  however,  be  admitted  that  the  evidence  in  regard  to  the 
structure  of  compounds  which  contain  more  than  one  element  of 
varying  valency  is  usually  far  from  conclusive.  Moreover,  elements 
in  some  cases  exert  valencies  different  from  those  deduced  from  the 
composition  of  their  hydrogen  or  halogen  derivatives,  e.g.  nitrogen 
in  NO,  chlorine  in  C1O2. 

The  maximum  valency  of  any  element  for  hydrogen  is  four,  but  in 
combination  with  the  halogens  much  higher  valencies  are  met  with, 
and  the  maximum  valencies  thus  obtained  usually  correspond  closely 
with  the  positions  of  the  respective  elements  in  the  periodic  table.1 
Thus,  with  scarcely  an  exception,  elements  of  the  fourth  group  form 
compounds  of  the  formula  EC14  [E  =  element],  those  of  the  fifth 
group  compounds  of  the  formula  EF5  or  EC15  (e.g.  PF6,  AsF6, 
SbF5,  TaCl5),  and  those  of  the  sixth  group  compounds  of  the  formula 
EF6  (e.g.  SF6,  SeF6,  TeF6,  WC16).  No  element  of  the  seventh 
group  forms  compounds  of  the  type  EFr,  and  the  only  element  of 
the  eighth  group,  the  octavalency  of  which  has  been  conclusively 
established,  is  osmium  from  the  compound  OsF8.  Of  the  compounds 
of  the  seventh  group,  the  chlorine  and  manganese  atoms  in  potassium 
perchlorate  and  permanganate  respectively  are  usually  regarded  as 

1  It  is,  of  course,  necessary  in  such  cases  to  show  that  the  simplest  formula 
represents  the  true  molecular  formula,  and  this  is  done  either  by  vapour  density 
measurements  or  by  determinations  of  the  molecular  weight  in  solution  (p.  195). 


5So     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

septavalent  (p.  550).  The  octavalency  of  ruthenium  may  be  inferred 
from  the  existence  of  the  compound  RuO4  (corresponding  with  OsO4X 
but  hitherto  no  evidence  has  been  obtained  that  the  other  metals  of 
the  eighth  group  can  exert  so  high  a  valency.  The  valency  also 
depends  on  external  conditions,  such  as  the  temperature  ;  it  tends  to 
diminish  as  the  temperature  rises.  Thus  ferric  chloride  decomposes 
into  ferrous  chloride  at  a  high  temperature,  the  valency  of  the  iron 
diminishing  from  three  to  two,  and  many  other  examples  of  the  same 
kind  can  be  given. 

.Theories  of  Valency — The  conception  of  valency  sketched 
above  has  rendered  very  valuable  service  in  Organic  Chemistry, 
mainly  because  carbon  and  hydrogen,  two  of  the  chief  elements 
present  in  organic  compounds,  show  a  constant  valency,  and  oxygen, 
also  present  in  a  great  number  of  organic  compounds,  is  almost  in- 
variably bivalent  In  Inorganic  Chemistry,  however,  the  application 
of  the  doctrine  of  valency  has  proved  much  less  useful,  mainly  on 
account  of  the  fact  that  the  same  element  may  show  different  valencies. 
The  so-called  "  molecular  compounds  "  (p.  272)  have  proved  particularly 
difficult  to  bring  into  line  with  the  conception  of  valency  in  the  above 
simple  form.  Thus  though  the  ordinary  atomic  valencies  both  in 
sodium  chloride  and  water  are  satisfied,  these  two  compounds  never- 
theless combine  to  form  a  definite  compound  of  the  formula  NaCI, 
2H2O  (p.  400).  Numerous  other  so-called  molecular  compounds  have 
been  mentioned  in  the  course  of  the  book. 

Although  many  theories  of  valency  have  been  put  forward,  none 
has  so  far  met  with  general  acceptance.  The  most  suggestive  and 
comprehensive  is  that  due  to  Werner,1  which  will  first  be  described. 

Werner's  Theory  of  Valency — The  theory  of  valency  due 
to  Werner  is  based,  in  the  first  instance,  on  the  investigation  of  the 
so-called  "molecular  compounds,"  including  compounds  of  salts 
with  water,  known  as  hydrates,  compounds  of  salts  with  ammonia, 
known  as  ammines,  e.g.  PtCl4,  2NH8,  and  complex  halogen  salts,  e.g. 
PtCl4,  2HC1.  It  was  at  first  thought  that  there  is  a  sharp  distinction 
between  such  compounds  and  those  in  which  the  atoms  can  be 
regarded  as  being  linked  together  by  the  ordinary  valencies,  but,  as 
we  shall  see  later,  no  such  sharp  distinction  exists. 

The  method  employed  in  investigating  the  constitution  of  such 
molecular  compounds  can  be  conveniently  illustrated  by  means  of 
the  compound  PtCl4,  2NH3.  This  compound  in  aqueous  solution 

1  Cf.  Neuere  Anschauungen  auf  dem  Gebeite  der  Anorganischen  Chemie,  2nd 
Edition,  1909.  English  translation  by  Hedley  (Longmans,  Green  &  Co.),  1911. 


MODERN  VIEWS  ON  VALENCY  581 

does  not  conduct  the  electric  current,  and  the  chlorine  atoms  must 
therefore  be  directly  associated  with  the  platinum — if  the  formula 
was  "of  the  type  PtNH3)Cl,  the  salt,  like  ammonium  chloride,  would 
contain  free  Cl'  ions.  As  regards  the  mode  in  which  the  NH8  groups 
are  bound,  the  following  are  the  possible  formulae  : — 

Cl 
Ck         /Cl-NHg          Ck          ,NH3)C1          Ck     |       NH3 

>P<  ;  >Pt  ;  ;>Pt: 

Cl/       \C1-NH3          Cl/          NH3)C1          Cl/   I    '--NH3 

Cl 

The  first  formula  is  excluded  because  it  is  possible  to  remove  two 
chlorine  atoms  by  reduction  without  affecting  the  NH3  groups.  The 
second  formula  is  excluded  for  the  reason  already  given,  that  the 
compound  behaves  quite  differently  to  ammonium  chloride,  and  there- 
fore the  chlorine  atoms  must  be  directly  attached  to  platinum.  That 
the  two  NH3  groups  are  independent  is  proved  by  the  possibility  of 

,  NH8 
converting  the  compound,  Cl4Pt  ,  the  formula  of  which  can  easily 

C1H 
,NH8 
be  established,  into  Cl4Pt  by  the  direct  action  of  ammonia.    By 

NH3 
similar  reasoning  it  can  be  shown  that  the  formula  of  "  chlorplatinic 

,C1H 
acid"  (PtCl4,2HCl)   should  be   represented   thus:   Cl4Pt  ,   and 

'••C1H 
OH2 
the  corresponding  dihydrate  as  Cl4Pt: 

OH2. 

The  platinum-ammonia  compound  just  described  is  one  of  a  series 
the  formula  of  the  members  of  which  l  (as  already  pointed  out)  are 
represented  on  Werner's  theory  as  follows  : — 

[Pt(NH8)6]""  C14  ;  [Pt(NH,)5Cl]-  C13  ;  [Pt(NH3)4Cl2]"  Cl,  ; 
[Pt(NH8)8ClJ-  Cl ;  Pt(NH3)2Cl4]  ;  [PtNH.ClJ  K  ;  [PtCl«r  K2. 
The  first  compound  has  a  high  electrical  conductivity  in  solution, 
and  all  the  chlorine  atoms  are  readily  precipitated  by  silver  nitrate. 
In  the  third  compound,  on  the  other  hand,  the  electrical  conductivity 
is  much  smaller,  and  only  two  of  the  chlorine  atoms  react  with  silver 
nitrate ;  finally,  as  already  mentioned,  the  fifth  compound  has 
practically  no  electrical  conductivity,  and  the  chlorine  does  not  react 
with  silver  nitrate.  Werner  accounts  for  this  behaviour  on  the  view 
1  The  second  compound  is  so  far  unknown,  but  its  formula  is  given  for  the  sake 
of  completeness. 


582     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

that,  as  already  shown  for  Cl4Pt(NH3)2,  the  groups  within  the  square 
brackets  are  directly  bound  to  the  platinum  in  the  inner  zone  ;  chlorine 
bound  in  this  way  behaves  like  chlorine  in  an  organic  compound,  and 
does  not  give  the  ordinary  reactions — in  other  words,  it  is  not  ionised. 
The  atoms  in  the  outer  zone,  on  the  other  hand,  are  less  intimately 
bound  to  the  central  atom  and  are  ionised.  The  number  of  groups 
attached  directly  to  the  central  atom  is  known  as  the  "  co-ordination 
number "  of  the  atom,  and  is  determined  from  the  behaviour  of  the 
compound,  as  in  the  example  quoted.  Examples  from  chromium, 
cobalt,  and  platinum  compounds  have  already  been  given. 

It  will  be  noticed  that  the  presence  of  ammonia  molecules  in  the 
inner  zone  does  not  alter  the  valency  of  the  central  atom.  Thus 
the  valency  of  the  complex  [Co(NH3)6]  is  three,  corresponding  with 
the  valency  of  the  simple  Co"*  ion ;  the  only  effect  of  the  addition 
of  the  ammonia  atoms  is  to  increase  the  stability  of  the  compounds 
containing  trivalent  cobalt  (p.  551).  The  presence  of  water  and 
of  certain  other  groups  is  also  without  effect  on  the  valency  of  the 
complex  cation.  The  NH3  groups  in  the  compound  [Co(NH3)6]Cl3 
can  be  gradually  displaced  by  H2O  molecules,  down  to  Co(H2O)6Cl3, 
the  complex  group  containing  the  central  atom  remaining  trivalent 
throughout.  On  the  other  hand,  for  every  NH3  molecule  displaced 
by  a  strongly  negative  univalent  atom  or  group,  the  positive  valency 
of  the  central  group  is  diminished  by  one  unit.  This  is  shown  very 
clearly  in  the  formulae  of  the  platinum  compounds  already  quoted. 
When  four  Cl  atoms  have  entered  the  inner  sphere,  the  four  possible 
valencies  of  the  platinum  are  neutralized  and  the  compound  is 
electrically  neutral ;  when  a  fifth  chlorine  atom  enters,  the  group 
has  become  negative  and  univalent,  with  two  Cl  atoms  it  is  negative 
and  divalent. 

So  far  we  have  not  discussed  the  nature  of  the  valencies  in  'such 
compounds,  but  this  is  a  matter  of  fundamental  importance.  Werner 
distinguishes  two  kinds  of  valencies — (i)  the  principal  valencies,  which 
are  measured  by  the  number  of  hydrogen  atoms  the  element  or  group 
can  combine  with  or  displace  ;  (2)  the  subsidiary  valencies,  which  bind 
groups  which  are  capable  of  existing  as  independent  molecules.  In 
writing  graphic  formulas  the  ordinary  valencies  are  represented  by 
continuous  lines,  the  subsidiary  valencies  by  dotted  lines.  It  is 
evident  that  the  principal  valencies  correspond  with  those  ordinarily 
accepted,  which  bind  atoms  or  groups  such  as  Cl,  NO2,  CH3.  The 
subsidiary  valencies,  on  the  other  hand,  bind  groups  such  as  water 
and  ammonia,  the  linking  of  which  cannot  be  represented  on  the 


MODERN  VIEWS  ON  VALENCY  583 

ordinary  theory  of  valency.  A  further  difference  between  the  two 
kinds  of  valency  is  that  some  at  least  of  the  groups  joined  to  the 
central  atom  by  principal  valencies  can  ionise,  whereas  groups  joined 
by  subsidiary  valencies  are  incapable  of  ionisation. 

It  should  be  stated,  however,  that  there  is  no  perfectly  sharp 
distinction  between  the  two  kinds  of  valency.  It  will  be  evident 
from  the  foregoing  that  the  co-ordinated  groups  may  be  joined 
either  by  principal  or  by  subsidiary  valencies ;  in  the  compound 
[Co(NH3)6]Cl3,  for  instance,  they  are  all  connected  to  the  central 
atom  by  subsidiary  valencies. 

It  is  a  very  remarkable  fact  that  the  co-ordination  number  is  six  for 
so  many  elements,  largely  independent  of  the  nature  of  the  co-ordinated 
groups,  and  Werner  is  of  opinion  that  it  is  more  a  question  of  space 
than  of  affinity  ;  six  is  as  a  rule  the  maximum  number  of  groups  that 
can  be  fitted  round  the  central  atom.  The  co-ordination  number  of 
carbon  is  4,  corresponding  with  its  ordinary  valency. 

The  valencies  of  an  atom  have  often  been  regarded  as  forces 
directed  from  certain  points  of  its  surface.  Werner  replaces  this 
view  of  separate  directed  forces  by  the  conception  of  affinity,  which 
is  regarded  as  an  attractive  force  acting  from  the  centre  of  the  atom 
uniformly  over  its  whole  surface.  When,  for  instance,  two  atoms 
such  as  Na  and  Cl  are  linked  together,  part  of  the  affinity  of  the 
sodium  atom  is  satisfied  by  the  chlorine  atom,  but  a  certain  amount 
of  "residual  affinity"  remains  over  on  each  atom  and  is  available 
for  linking  up  with  other  atoms.  This  conception  of  residual  affinity 
goes  back  to  the  time  of  Berzelius,  but  has  been  put  in  the  most 
definite  form  by  Werner.  The  amount  of  affinity  "  exchange  "  between 
two  atoms  depends  greatly  on  their  nature.  Thus  if  four  similar 
atoms  are  attached  to  carbon,  they  will  each  take  up  an  equal 
amount  of  affinity  and  will  arrange  themselves  so  that  the  greatest 
possible  saturation  of  affinities  takes  place.  A  little  consideration 
will  show  that  this  is  secured  by  the  atoms  occupying  the  four 
corners  of  a  tetrahedron,  at  the  centre  of  which  is  the  carbon  atom. 
When  the  four  atoms  or  groups  are  not  identical,  they  will  take  up 
different  amounts  of  affinity  from  the  carbon  atom  and  the  arrange- 
ment in  space  v/ill  no  longer  be  symmetrical.  These  considerations 
help  to  elucidate  the  view  of  Werner  already  referred  to,  that  the 
number  of  atoms  or  groups  bound  by  a  central  atom  is  determined 
more  by  space  considerations  than  by  affinity,  and  in  any  case 
cannot  be  accounted  for  on  the  hypothesis  of  separate  directed 
valencies.  It  should  be  emphasized  that  on  no  theory  of  valency  is 


584    A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 


there   any   proportionality   between    the    chemical    affinity  and    the 
number  of  "  valency  bonds." 

We  will  conclude  this  brief  account  of  Werner's  theory  by  reference 
to  the  interesting  cases  of  isomerism  met  with  in  the  case  of  the 
compounds  [Pt(NH3)2Cl2]  and  [Pt(NH3)2Cl4].  Each  of  these  com- 
pounds exists  in  two  entirely  distinct  forms  which  differ  in  many  of 
their  reactions.  Considering  first  the  compound  Pt(NH3)2Cl2,  the 
isomerism  cannot  easily  be  accounted  for  on  the  assumption  that 
the  groups  are  symmetrically  arranged  in  space  round  the  Pt  atom 
at  the  four  corners  of  a  tetrahedron,  as  in  such  a  case  the  groups 
could  be  interchanged  without  altering  their  mutual  relations,  and 
therefore  only  one  isomeride  is  to  be  expected.  If,  however,  the 
four  groups  are  in  the  same  plane  as  the  platinum  atom,  they  can 
be  arranged  in  the  two  alternative  ways  : 


Cl 


NH. 


NHc 


NH: 


(trans-position). 

k       /NH3 
^Pt/ 


Cl 


Cl  NH3 

(cis-position). 

or,  more  simply 

Cl\        /NH3 
\P|-/ 

Cl/     ^NHg  NH3/ 

On  the  other  hand,  the  two  isomers  of  the  formula  [Pt(NH3)2Cl4] 
can  be  represented  graphically  as  follows,  the  six  groups  being 
arranged  in  space  round  the  platinum  atom  at  the  corners  of  an 
octahedron. 

Cl  NH3 


NH 


Cl 


Pt 


Pt 


Cl 


NH3 


Cl 


NH, 


MODERN  VIEWS  ON  VALENCY  585 

Many  cases  of  isomerism  are  also  met  with  among  chromium  and 
cobalt  salts.  The  most  recent  investigations  lend  support  to  the  view 
that  there  is  no  essential  difference  between  valency  compounds  and 
molecular  compounds. 

Ab©gg's  Theory  of  Valency J — Abegg's  theory  of  valency 
may  be  regarded  as  a  development  of  the  views  of  Berzelius  (p.  278) 
on  the  electric  polar  character  of  chemical  compounds,  and  is  not  in 
opposition  to  Werner's  theory,  though  in  some  respects  more  definite 
than  the  latter.  He  assumes  that  the  valencies  of  the  elements  are 
polar  in  character,  that  every  element  possesses  a  positive  as  well  as 
a  negative  maximum  valency,  and  that  the  sum  of  the  two  maximum 
valencies  as  for  all  elements  equal  to  8.  The  stronger  valencies  of 
each  element,  which  are  also  less  in  number  «4)  are  termed  the 
normal  valencies  ;  the  other  valencies  of  opposite  polarity  are  termed 
contra  valencies.  In  the  case  of  the  metals  the  normal  valencies  are 
positive  and  the  contra  valencies  negative  ;  in  the  case  of  the  non- 
metals,  the  normal  valencies  are  negative  and  the  contra  valencies  are 
positive.  The  positive  maximum  valency  corresponds  therefore  with 
the  group  number  in  the  periodic  system — in  other  words,  with  the 
generally  accepted  maximum  valencies  of  the  elements  (p.  381).  The 
arrangement  for  one  short  period  is  as  follows  : — 

Group       ...        I  II  in  iv  v  vi  vn 

Example  .         .        .  Na  Mg  Al  Si  P        S  Cl 

Normal  Valencies     .      +i  +2  +3  +4  —3  —2  —  I 

Contra  Valencies      .  (-7)  (-6)  (-5)  -4  +5  +6  +7 

The  smaller  the  number  of  normal  valencies  the  more  pronounced 
is  the  polar  character,  e.g.  Na  and  Cl ;  on  the  other  hand,  for  elements 
near  the  middle  of  the  group  the  number  of  normal  and  contra 
valencies  approximate  more  and  more  in  number  and  strength.  This, 
according  to  Abegg,  is  the  reason  why  carbon  unites  readily  both  with 
positive  (H,  Zn,  etc.)  and  also  negative  atoms  (O,  Cl),  and  also  why 
carbon  atoms  unite  so  readily  with  each  other.  The  formation  of 
molecular  compounds  is  ascribed  to  the  participation  of  contra 
valencies,  which  are  latent  under  ordinary  circumstances  in  such 
compounds  as  water  and  ammonia. 

In  some  respects  Abegg's  theory  affords  a  satisfactory  explanation 
of  the  facts,  e.g.  the  fact  that  the  valency  of  chlorine  to  positive 
hydrogen  is  i,  and  to  negative  oxygen  7  (in  C12O7),  and  also  that 
1  Zeitsch.  Anorg.  CAem.,  1904,  jp,  330. 


586     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

sulphur  is  divalent  to  hydrogen  and  sexavalent  to  electro-negative 
fluorine.  There  is  no  evidence,  however,  that  the  alkali  metals  can 
meet  a  maximum  negative  valency  of  7,  as  the  theory  requires  : 
potassium  salts,  for  example,  show  very  little  tendency  to  form  mole- 
cular compounds. 

Moreover,  as  Kauffmann  has  pointed  out,  the  formation  of  ethane, 
H3C-CH3,  by  the  action  of  metallic  sodium  on  CH3Br  (p.  346) 
is  difficult  to  reconcile  with  Abegg's  theory.  The  residue  CH3  must 
be  positive,  as  it  is  derived  from  a  compound  with  electro-negative 
Br,  but  the  compound  H3C-CH3,  resulting  from  the  association  of  two 
such  groups,  is  not  polar  in  character.  The  conception  of  polar 
valencies  has  hitherto  failed  to  justify  itself  as  far  as  organic  com- 
pounds are  concerned. 

It  should  be  mentioned  that  views  very  similar  to  those  of  Abegg 
were  previously  put  forward  by  Mendeleeff. 

J.  J.  Thomson's  explanation  of  valency  on  the  basis  of  his  electronic 
theory  of  the  constitution  of  the  atom  will  be  briefly  referred  to  in  the 
next  chapter  on  Radio-activity. 


CHAPTER  XXXVII 
RADIO-ACTIVITY 

ONE  of  the  most  striking  properties  of  the  Rontgen  or  X-rays  is 
the  fluorescence  to  which  they  give  rise  on  screens  of  certain 
materials.  In  the  course  of  an  examination  of  a  number  of  phos- 
phorescent and  fluorescent  substances,  undertaken  in  order  to  find 
out  if  they  give  rise  to  radiation  of  any  kind,  Becquerel,  in  1896, 
discovered  that  uranium  salts  affected  a  photographic  plate  through 
several  layers  of  black  paper.  This  effect  was  clearly  due  to  radiation 
from  the  uranium.  The  Becquerel  rays,  as  this  type  of  radiation 
came  to  be  called,  had  the  further  property  of  rendering  the  air  in 
their  neighbourhood  a  conductor  of  electricity  and  therefore  of  dis- 
charging a  gold  leaf  electroscope.1  Substances  giving  out  rays  of 
this  type  are  said  to  be  radio-active. 

A  little  later,  M.  and  Mme.  Curie  found  that  several  minerals 
containing  uranium,  notably  pitchblende  from  Austria,  were  found  to 
be  considerably  more  radio-active  than  pure  uranium  compounds 
or  even  than  uranium  itself.  It  therefore  appeared  probable  that 
the  radio-activity  was  not  due  to  uranium,  but  to  some  other  con- 
stituent in  uranium  ores,  and  M.  and  Mme  Curie  set  themselves  the 
task  of  isolating  the  substance.  Ultimately,  two  very  active  substances 
were  obtained  ;  one,  closely  allied  in  its  chemical  characters  to 
bismuth,  was  named  polonium  (from  Poland,  Mme.  Curie's  native 
country),  the  other,  closely  allied  chemically  with  barium,  was  termed 
radium.  The  amount  of  radium  present  in  an  average  sample  of 
pitchblende  is  extremely  small,  not  exceeding  i  part  in  10  million 
parts  of  the  ore,  and  polonium  occurs  in  still  smaller  proportion  (not 
exceeding  i  part  in  10,000  million  parts  of  pitchblende).  Pure 
radium  salts  are  about  2  million  times  more  radio-active  than  uranium 
compounds. 

1  A  gold  leaf  electroscope  consists  essentially  of  two  insulated  gold  leaves  (sus- 
pended close  together),  which  diverge  when  electrically  charged,  but  fall  together 
when  the  air  in  contact  with  them  becomes  conducting. 

587 


588     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

The  possibility  of  isolating  elements  present  in  such  small  pro- 
portion is  due  to  the  great  readiness  with  which  radio-activity  can 
be  detected.  Up  to  that  time  the  use  of  the  spectroscope  was  the 
most  sensitive  means  of  finding  minute  quantities  of  substances, 
but  those  substances  which  are  radio-active  can  be  detected  in  much 
more  minute  quantities.  The  presence  of  as  little  as  io~12  gram  of 
radium  can  be  proved  in  this  way. 

At  a  later  period,  a  third  radio-active  substance,  called  actinium, 
was  found  in  pitchblende.  Neither  polonium  nor  actinium  appears 
to  have  been  obtained  pure.  Early  in  her  investigations,  Mine.  Curie 
discovered  that  thorium  compounds  also  show  radio-activity. 

Properties  of  Radium — Radium  is  a  well-defined  element,  of 
atomic  weight  226,  and  it  occupies  a  position  in  the  periodic  table 
as  the  last  member  of  the  barium  group.  All  the  work  on  radio- 
activity has  hitherto  been  done  with  radium  chloride,  RaCl2,  and 
radium  bromide,  RaBr2,  but  quite  recently  (1910)  the  element  itself 
was  isolated  by  Mme.  Curie  and  Debierne.  They  subjected  a 
solution  of  radium  chloride  to  electrolysis  with  a  mercury  cathode, 
distilled  off  the  mercury,  and  obtained  radium  as  a  brilliant  white 
metal,  which  melted  sharply  at  700°  and  volatilized  slightly  at  the 
same  temperature.  Like  the  other  metals  of  the  alkaline  earths, 
radium  is  readily  attacked  by  water,  hydrogen  being  evolved,  and  it 
tarnishes  rapidly  in  the  air. 

Investigation  of  the  radiation  given  off  by  radium  (and  by  uranium 
ores)  has  shown  that  three  distinct  kinds  of  rays,  known  respectively 
as  a,  |3,  and  y  rays,  can  be  distinguished.  All  of  them  affect  a 
photographic  plate,  discharge  a  gold  leaf  electroscope,  and  render 
materials  such  as  zinc  sulphide  and  zinc  silicate  (willemite)  fluorescent, 
but  their  relative  activities  are  very  different. 

The  a  rays  have  very  little  penetrating  power,  being  completely 
stopped  by  three  or  four  layers  of  thin  aluminium  foil,  by  a  sheet  of 
paper,  or  by  passing  through  a  layer  of  air  2  to  3  cm.  thick.  Their 
electric  effect  is  very  marked.  They  are  slightly  deflected  by  a 
magnet,  the  direction  of  the  deflection  showing  that  they  are  positively 
charged.  It  is  now  generally  accepted  that  when  an  a  particle  gives 
up  its  charges,  an  atom  of  helium  remains. 

The  /3  rays  have  at  least  100  times  the  penetrating  power  of  the  a 
rays.  Rutherford  showed  that  on  passing  through  100  thin  plates 
of  aluminium  their  intensity  was  only  reduced  by  half.  They  are 
readily  deflected  by  a  magnet,  the  direction  showing  that  they  are 
negatively  charged,  and  their  mass  is  only  about  j^  of  that  of  th* 


RADIO-ACTIVITY  589 

hydrogen  atom.      It   is  generally  agreed  that  the  $   particles  are 
negative  electrons — actual  units  of  electricity. 

The  y  rays  have  a  penetrating  power  much  greater  than  the  /3  rays 
and  affect  a  photographic  plate  strongly.  They  travel  with  the 
velocity  of  light.  They  appear  to  be  identical  with  Rontgen  or 
X-rays,  and  are  therefore  not  particles,  but  a  wave  motion  in  the 
ether. 

Besides  these  rapidly-moving  particles,  radium  compounds  con- 
tinuously give  off  a  heavy  vapour  known  as  radium  emanation. 
Thorium  compounds  give  rise  to  an  analogous  emanation.  The 
emanations  are,  however,  unstable,  and  are  only  temporarily  radio- 
active. Radium  emanation  at  first  gives  off  only  a  rays.  At  a  later 
stage  it  gives  of  /3  and  y  rays  and  other  substances  are  also  formed. 
It  is  usual  to  express  the  stability  of  a  radio-active  substance  in 
terms  of  the  time  taken  for  its  activity  to  fall  to  half  its  original 
value.  The  half-time  period  for  radium  emanation  is  about  3.85 
days. 

Chemically,  radium  emanation  behaves  like  a  gas  of  the  helium 
series  (p.  209),  and  in  a  vacuum  tube  shows  a  characteristic  spectrum 
of  bright  lines.  It  is  unaffected  by  passing  through  a  hot  tube  or 
through  acids  ;  it  is  condensed  by  passing  through  a  tube  immersed 
in  liquid  air,  but  vaporizes  again  when  the  temperature  is  allowed 
to  rise.  The  liquid  emanation  boils  at  -  62°  C.  It  is  colourless  when 
first  condensed ;  when  the  temperature  is  lowered  it  freezes,  and 
at  the  temperature  of  liquid  air  glows  with  a  bright  rose  colour. 
The  density  of  the  emanation  in  the  form  of  vapour  has  been 
determined  from  its  rate  of  diffusion,  and  quite  recently  has  been 
determined  directly.  The  value  obtained  is  about  220,  which  would 
bring  it  into  the  helium  series,  below  xenon. 

The  most  surprising  fact  connected  with  the  emanation  is  that  one 
of  the  products  of  its  decay  is  helium  (Ramsay  and  Soddy).  When 
a  little  of  the  emanation  is  collected  in  a  tube  and  its  spectrum 
examined  at  intervals,  no  helium  lines  are  noticeable  at  first,  but 
after  a  time  they  appear  and  gradually  become  more  intense.  In 
the  light  of  what  has  been  mentioned  as  to  the  connexion  between 
a  particles  and  helium,  the  appearance  of  the  latter  element  will  be 
understood.  While  the  emanation  is  breaking  up  it  gives  off  a 
particles,  and  when  these  lose  their  charges  they  become  ordinary 
helium  atoms. 

In  the  process  of  decay  of  the  emanation,  six  further  substances 
appear  to  be  formed  in  succession,  known  respectively  as  radium 


590     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

A,  B,  C,  D,  E,  and  F.  The  main  part  of  radium  Cx  changes  to 
radium  D,  but  a  small  amount  of  radium  C2  is  formed  as  a  side 
product. 

Disintegration  Theory  of  Radio-activity—  We  have  now 

to  consider  whether  any  theory  can  be  suggested  to  account  for  the 
remarkable  facts  stated  in  the  previous  section.  In  this  connexion 
two  further  important  points  should  be  mentioned  :  (i)  radium  com- 
pounds are  continually  giving  out  heat  and  maintain  themselves  at 
a  temperature  of  i°  to  2°  above  their  surroundings  ;  (2)  the  radio- 
activity of  a  series  of  compounds  of  a  radio-active  element  is  directly 
proportional  to  the  amount  of  the  element  present.  As  has  already 
been  pointed  out,  radium  is  undoubtedly  an  element,  with  a  definite 
spectrum  and  a  definite  place  in  the  periodic  system.  The  simplest 
explanation  of  all  the  facts  we  have  mentioned  is  that  atoms  of  this 
element  are  continually  disintegrating  or  exploding,  throwing  off 
charged  particles  and  forming  substances  of  smaller  atomic  weight. 
The  first  stage  consists  in  the  expulsion  of  an  a  particle  (helium  atom) 
and  the  formation  of  radium  emanation,  which  presumably  has  an 
atomic  weight  4  units  less  than  that  of  the  radium  atom.  The  atoms 
of  the  emanation  in  their  turn  explode,  the  first  stage  being  the 
expulsion  of  an  a  particle  with  formation  of  radium  A.  Radium  A 
in  its  turn  expels  an  a  particle  and  radium  B  results.  The  latter 
in  its  turn  breaks  down  into  radium  C15  in  which  transformation 
/3  particles  are  expelled,  and  so  on.  The  time  occupied  in  these 
different  stages  is  very  different.  In  the  case  of  radium,  the  pro- 
portion of  atoms  exploding  is  very  small,  and  it  has  been  calculated 
that  it  requires  abcut  2000  years  for  the  activity  of  radium  to  fall  to 
half  its  value  On  the  other  hand,  the  half-time  period  for  the 
emanation  is  about  3.85  days,  and  for  radium  A  to  radium  B  only 
three  minutes.  The  successive  stages  in  the  disintegration  of  radium 
are  given  in  the  accompanying  table  ;  the  rays  given  off  (if  any) 
at  each  stage  are  given  in  brackets,  and  the  half-time  period  below  : 


Radium  (*)  ->  Emanation  (*)  ->  A(«)  ->  B(/3,y)  ->  Cj^./S.y)  ->  D  ->  E(/3,y)  ->  F(«) 
2000  years  3.85  days  3  min.  27  min.  19.5  min.  16.5  yrs.  5  dys.  136  dys. 

Radium  F  is  polonium,  discovered  in  pitchblende  by  Mme.  Curie. 
It  breaks  down  much  more  rapidly  than  radium,  but  the  final  product 
is  unknown.  It  is  doubtless  a  non-radio-active  and  stable  element, 
and  may  possibly  be  lead  (atomic  weight  207),  which  is  present  in 
all  uranium  minerals  (cf.  p.  594)- 


RADIO-ACTIVITY  591 

On  the  disintegration  theory  the  source  of  the  energy  given  out 
in  a  radio-active  change  is  the  internal  energy  of  the  atom  ;  the 
process  of  breaking  up  results  in  the  formation  of  a  succession  of 
substances  with  less  and  less  energy.  Curie  and  Laborde  found 
that  I  gram  of  radium  evolves  100  calories  per  hour.  It  has  been 
calculated  that  the  heat  given  out  in  the  complete  disintegration  of 
i  gram  of  radium  would  raise  100  tons  of  water  i°  in  temperature  ; 
or,  in  mechanical  units,  the  available  energy  in  I  gram  of  radium 
would  suffice  to  raise  400  tons  I  mile  high.  An  alternative  illustration 
of  the  enormous  amounts  of  energy  concerned  is  that  in  the  complete 
change  of  1.3  cubic  millimetres  of  emanation  10,540  calories  are 
evolved — a  quantity  of  energy  four  million  times  greater  than  is 
evolved  by  the  same  volume  of  hydrogen  and  oxygen  when  they 
explode  to  form  water.  This  gives  us  some  idea  of  the  immense 
stores  of  energy  in  the  atom. 

The  chemical  activities  of  radium,  for  example  the  decomposition  of 
water  into  hydrogen  and  oxygen,  change  of  oxygen  to  ozone,  etc.,  are 
doubtless  connected  with  the  liberation  of  this  large  amount  of  energy. 

Radio-active  change  is  entirely  spontaneous,  and  no  method  of 
initiating  it  or  influencing  its  rate  is  known.  In  particular,  the  widest 
range  of  temperature,  from  that  of  liquid  hydrogen  to  2000°,  does  not 
appear  to  affect  the  speed  of  atomic  disintegration  in  the  smallest 
degree. 

Radio-activity  of  Uranium  and  the  Origin  of  Radium 
— As  radium  is  continuously  breaking  down,  it  is  evident  that  its 
presence  in  pitchblende  can  only  be  accounted  for  on  the  view  that 
it  is  continuously  being  re-formed  from  something  else.  As  the 
quantity  of  radium  in  uranium  ores  is  approximately  proportional  to 
the  quantity  of  uranium  present,  it  was  natural  to  suppose  that  the 
latter  is  the  parent  of  radium.  If  this  were  the  case,  radium  should 
be  produced  at  an  easily  measurable  rate  in  uranium  solutions. 
Experiments  made  to  test  this  point  showed  that  radium,  if  pro- 
duced at  all,  was  produced  much  less  rapidly  than  would  be  the  case 
if  it  were  obtained  directly  from  uranium  or  its  first  disintegration 
product,  uranium  X.  This  difficulty  was  overcome  by  the  discovery 
of  ionium  (Boltwood,  1907),  an  intermediate  substance  of  slow  rate  of 
transformation  between  uranium  and  radium.  The  uranium  series  is 
as  follows,  the  half-time  periods  being  given  in  the  second  line  : — 

Uranium  (a)  -»  UrX(/3-f  y)  ->  Ionium  (a)  ->  Radium  (a) 
8  x  io9  years          24.6  days         200,000  years      2000  years. 


592     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

As  uranium  in  disintegrating  expels  two  a  particles  per  atom,  it  is 
possible  that  it  is  not  a  single  element  but  a  mixture  of  two,  chemically 
non-separable,  differing  in  atomic  weight  by  four  units,  and  both 
expelling  a  rays.  Although  the  amount  of  ionium  present  in  uranium 
minerals  is  proportional  to  that  of  the  uranium,  and  therefore  it 
appears  probable  that  ionium  is  produced  by  the  disintegration  of 
uranium,  it  should  be  mentioned  that  the  direct  production  of  ionium 
from  pure  uranium  or  uranium  X  has  not  hitherto  been  observed. 
Ionium  is  chemically  non-separable  from  thorium. 

The  Actinium  Series  —  The  discovery  of  actinium  in 
pitchblende  has  already  been  referred  to.  It  has  been  shown  by 
Boltwood  that  the  amount  of  actinium  present  in  uranium  minerals 
is  approximately  proportional  to  the  percentage  of  radium.  This 
probably  indicates  that  actinium,  like  radium,  must  belong  to  the 
uranium  series.  For  reasons  which  cannot  be  considered  here,  it 
seems  unlikely  that  actinium  belongs  to  the  main  uranium-radium 
series  already  described.  It  may  be  a  branch  product  of  this  series, 
one  of  the  members  disintegrating  in  two  ways,  but  the  evidence  so 
far  available  is  not  conclusive  on  this  point.  The  determination  of 
the  atomic  weight  of  actinium  would  afford  valuable  information  on 
these  questions.  The  actinium  series  with  the  half-value  periods  is 
as  follows  : — 

Actinium-^  Radio- Act. (<*,£)-»  Act.X(a)->  Emanation  (a)-» 

?                   19.5  days.  10.2  days.  3.9  sees. 

Act.A(a)-»                B(/3)->  C(a)->  D(/3,y) 

0.002  sec.             36  mins.  2.1  mins.  4.71  mins. 

The  Thorium  Series  —  The  radio-activity  of  thorium 
compounds  was  discovered  simultaneously  and  independently  by 
G.  C.  Schmidt  and  Mme.  Curie  (1898).  The  thorium  series  is  a 
lengthy  one.  One  of  the  most  important  members  is  thorium  X,  the 
careful  investigation  of  which,  by  Rutherford  and  Soddy,  led  these 
investigators  to  put  forward  the  disintegration  theory.  They  dis- 
covered that  when  ammonia  was  added  to  a  thorium  solution,  the 
thorium  was  completely  precipitated,  but  the  filtrate,  free  from  thorium, 
was  found  to  contain  a  large  part  of  the  activity.  On  evaporation,  a 
small  residue — thorium  X — was  obtained  which,  weight  for  weight,  was 
several  thousand  times  more  active  than  the  thorium  from  which  it  was 
obtained.  On  examining  the  products  a  month  later,  it  was  found 
that  thorium  X  was  no  longer  active,  whilst  the  thorium  had  completely 
regained  its  activity.  The  explanation  in  terms  of  the  disintegration 


RADIO-ACTIVITY  593 

theory  will  be  evident.     The  thorium  series,  with  the  half-  value  periods, 
is  as  follows  :  — 

Thorium  (&)->  Mesoth.  i->  Mesoth.  2(/3,y)->  Radioth.(a)->  Th.X(a,/3->) 
i.3Xio10yrs.     5.5  yrs.  6.2  hrs.  2  yrs.  3.65  days. 


54  sees.        0.14  sec.  10.6  hrs.  60  mins.  3.1  mins. 

The  final  product,  produced  by  the  break-up  of  thorium  D,  is 
doubtless  a  non-radio-active  and  stable  element,  but  its  identity  has 
not  been  established.  The  final  member  of  the  actinium  series  is 
also  unknown.  On  the  basis  of  a  general  principle  discussed  in  the 
following  section,  it  appears  that  the  unknown  end-products  of  the 
three  disintegration  series  belong  to  the  same  position  in  the  periodic 
table,  namely,  that  occupied  by  lead. 

The  Radio  -  Elements  and  the  Periodic  Law, 
Isotopic  Elements1  —  It  has  already  been  mentioned  (p.  592) 
that  ionium  is  identical  in  chemical  behaviour  with  thorium, 
and  therefore  these  two  elements  cannot  be  separated  by  chemical 
means.  The  analogy  between  the  two  elements  is  so  far-reaching 
that  even  their  spectra  are  believed  to  be  identical.  Further  in- 
vestigation has  shown  many  other  instances  of  this  remarkable 
phenomenon.  Thus  radium,  mesothorium  I,  thorium  X,  and  actinium 
X  are  chemically  non-separable,  as  are  lead,  radium  B,  thorium  B, 
actinium  B,  and  radium  D.  These  results  have  within  the  last  two  or 
three  years  led  to  one  of  the  most  remarkable  generalizations  in 
chemistry,  and  have  thrown  a  flood  of  light  on  the  nature  of  the 
periodic  system.  Different  elements  occupying  the  same  place  in 
the  periodic  table  are  said  to  be  isotopic. 

In  order  to  understand  these  results,  it  is  necessary  to  consider 
a  little  more  fully  the  effect  of  the  expulsion  of  a  and  ft  rays  from  a 
disintegrating  element.  The  a  rays,  as  already  mentioned,  consist  of 
an  atom  of  helium  associated  with  two  positive  charges.  The  loss  of 
two  positive  charges  from  an  atom  affects  the  valency  of  the  product 
exactly  as  an  ordinary  electro  -chemical  change  of  valency.  For 
example,  if  the  atom  were  initially  in  Group  IV.,  its  ion  is  quadrivalent 
and  carries  four  atomic  charges  of  positive  electricity.  The  expulsion 
of  an  a  particle  with  its  two  charges  brings  it  into  the  divalent  Group 
II  ,  non-separable  from  radium.  On  the  other  hand,  the  loss  of  a  ft 

1  Compare  Soddy,  The  Chemistry  of  the  Radio-  Elements,  1913,  Part  II. 
(Longmans). 

38 


594     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

particle,  which  is  a  negative  electron,  increases  the  positive  valency 
of  the  product  by  one.  Radium  B,  for  example,  isotopic  with  lead, 
expels  a  /3  particle  and  becomes  radium  C,  isotopic  with  bismuth,  the 
mass  being  in  this  case  practically  unchanged.  When  one  a  and  two 
P  particles  are  expelled,  the  product  occupies  the  same  position  in 
the  periodic  table  as  the  original  element,  and  the  two  have  identical 
chemical  properties,  although  of  different  mass.  It  should  be  pointed 
out  that  isotopes  are  not  confined  to  the  members  of  one  disintegration 
series.  Whenever  two  or  more  elements  fall  into  the  same  position  in 
the  periodic  table,  they  are  chemically  non-separable  and  identical, 
quite  independently  of  their  origin  or  of  their  atomic  mass. 

Since  this  generalization  was  brought  forward,  its  validity  has  been 
tested  in  two  ways.  In  the  first  place  it  was  used  to  predict  the 
chemical  behaviour  of  radio- elements  which  had  not  then  been 
adequately  examined.  Thus  it  was  suggested  that  actinium  D, 
thorium  D,  and  radium  C2  should  be  isotopic  with  thallium,  and  this 
prediction  has  already  been  verified  by  Fleck  for  actinium  D  and 
thorium  D.  It  should  be  mentioned,  however,  that  thallium  chloro- 
platinate  is  less  soluble  than  the  chloroplatinates  of  the  other  two 
elements,  whilst  in  their  other  reactions  actinium  D  and  thorium  D 
resemble  thallium  completely.  This  shows  that  though  isotopes 
resemble  each  other  so  closely  that,  with  the*  exception  mentioned, 
they  have  so  far  proved  chemically  non-separable,  the  discovery  of 
small  differences  in  behaviour  which  may  be  utilized  to  separate 
them  is  by  no  means  improbable. 

A  second  highly  important  consequence  of  the  theory  is  that  a 
"chemical  element  "is  not  necessarily  homogeneous,  and  its  atomic 
weight  may  be  a  mean  value  rather  than  a  natural  constant.  More- 
over, the  atomic  weights  of  elements  from  different  sources  will  not 
necessarily  be  identical,  as  they  may  be  expected  to  contain  the  constit- 
uents in  different  proportions.  This  consequence  of  the  theory  was 
first  tested  in  the  case  of  lead  by  Soddy  and  Hyman.  According  to 
the  theory)  the  unknown  end-products  of  all  the  known  disintegration 
series  fall  into  the  place  in  the  periodic  table  occupied  by  lead.  The 
calculated  atomic  weight  of  the  isotope  derived  from  uranium  is  206, 
and  that  of  the  isotope  derived  from  thorium  is  208.4,  since  they  are 
formed  from  radium  226.0,  and  thorium  232.4,  by  the  loss  of  five  and 
six  atoms  of  helium  respectively  (Fajans,  Soddy).  The  mineral 
thorite  proved  particularly  suitable  for  the  investigation,  as  it  consists 
mainly  of  thorium  and  uranium.  The  lead  in  the  mineral  proved  to 
have  an  atomic  weight  of  208.4,  whilst  the  accepted  atomic  weight  of 


RADIO-ACTIVITY  595 

lead  is  207.1.  The  conclusion  that  the  atomic  weight  of  lead  differs 
according  to  its  origin  was  fully  confirmed  later  by  Richards  and 
Lembert  (1914),  who  obtained  the  following  results :  lead  from 
uraninite  (North  Carolina),  206.40  ;  from  pitchblende  (Joachimsthal), 
206.57  ;  from  carnotite  (Colorado),  206.59  ;  from  Ceylon  thorianite, 
206.82  ;  from  English  pitchblende,  206.86  ;  ordinary  lead,  207.15. 

The  Structure  of  the  Atom.  J.  J.  Thomson's  Theory 
of  Valency l — Although  many  attempts  have  been  made  in  recent 
years  to  account  for  chemical  phenomena  on  the  basis  of  the  electron 
theory,  so  far  none  of  the  attempts  has  met  with  general  acceptance. 
It  will  therefore  be  sufficient  for  our  present  purpose  to  give  a  brief 
account  of  the  most  interesting  and  suggestive  of  these  efforts,  that 
due  to  J.  J.  Thomson.  It  has  been  shown  by  experiments  on  the 
scattering  of  X  rays  that  the  number  of  electrons  in  an  atom  is  about 
half  the  atomic  weight,  expressed  in  the  ordinary  units.  Thus  the 
hydrogen  atom  contains  only  one  electron,  and  the  thorium  atom 
over  100.  As  the  electrons  carry  a  large  negative  charge"  and  the 
atom  is  electrically  neutral,  it  must  be  assumed  that  the  atom  con- 
tains also  an  exactly  equivalent  amount  of  positive  electricity. 
Thomson  assumes  that  the  positive  electricity  is  uniformly  distributed 
within  the  circumference  of  the  atom,  and  that  the  electrons  are  in 
continuous  rotation  in  concentric  rings  round  the  centre  of  the  atom. 
The  electrons  therefore  tend  to  be  driven  outwards  by  centrifugal 
force,  and  are  also  acted  on  by  mutual  repulsive  forces  and  by  the 
attraction  of  the  total  positive  charge.  The  question  of  the  stability 
of  such  systems  is  then  considered.  It  is  calculated  that  when  all  the 
corpuscles  are  moving  in  one  circle  five  is  the  highest  number  that 
can  be  present.  Therefore  when  the  number  of  electrons  is  great, 
they  must  be  arranged  in  a  number  of  concentric  circles,  e.g.  when 
the  total  number  is  60,  they  are  moving  in  five  concentric  circles,  20 
in  the  outermost  ring,  16  in  the  next,  13  in  the  next,  8  in  the  fourth, 
and  3  in  the  innermost  circle.  The  stability  of  the  outer  ring 
increases  with  the  number  of  electrons  it  contains.  If  it  is  not  very 
stable  it  can  lose  an  electron  under  the  influence  even  of  weak  ex- 
ternal forces  and  an  atom  with  unit  positive  charge,  that  is,  a  univalent 
ion,  such  as  hydrogen  or  potassium,  is  obtained.  By  the  loss  of  two 
electrons  a  positively  charged  divalent  ion  is  obtained,  and  so  on.  The 
hypothesis  thus  affords  a  simple  explanation  of  the  variable  valency 
of  an  atom.  Similar  arrangements  of  corpuscles  occur  periodically  as 

1  J.  J.  Thomson,  The  Corpuscular  Theory  of  Matter,  1907  (Constable),;  The 
Atomic  Theory,  1914  (Clarendon  Press) ;  Phil.  Mag.,  1914,  27,  780. 


596     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

the  number  of  electrons  in  the  atom  increases.  Thus  the  sets  5,  i  ;  n, 
5,  i  ;  15,  n,  5,  i,  and  17,  15,  11,  5,  i  may  be  expected  to  have  much  in 
common,  and  this  would  be  in  harmony  with  the  periodic  /recurrence 
of  physical  and  chemical  properties  with  increasing  atomic  weight.  In 
his  most  recent  papers  the  author  distinguishes  between  fixed  corpuscles, 
which  are  incapable  of  exerting  any  great  attraction  on  neighbouring 
atoms,  and  mobile  corpuscles,  near  the  surface  of  the  atom,  which  have 
to  be  "  fixed  "  in  order  that  the  atom  may  be  "  saturated."  The  number 
of  these  mobile  corpuscles  in  an  atom  of  an  element  is  equal  to  the 
number  of  the  group  in  which  the  element  is  placed  according  to 
the  periodic  system.  » 

Rutherford  considers  that  the  scattering  of  a  particles  in  their 
passage  through  matter  cannot  be  satisfactorily  accounted  for  on  the 
assumption  that  positive  electricity  is  distributed  throughout  a  sphere 
of  radius  comparable  with  that  of  the  atom,  and  suggests  1  that  the 
atom  consists  of  a  positive  charge  concentrated  at  its  centre  and 
surrounded  by  a  distribution  of  electrons  to  render  it  electrically 
neutral. 

It  is  generally  supposed  that  chemical  and  electro-chemical  proper- 
ties are  controlled  by  the  outer  ring  of  electrons,  and  that  radio-active 
change  is  concerned  with  the  nucleus. 

Atomic  Numbers — When  cathode  rays  in  an  evacuated  tube  impinge  on 
a  plate  of  metal  or  other  element  (the  anti-cathode)  Rb'ntgen  or  X-rays  of  two 
definite  wave-lengths  are  obtained,  these  X-ray  spectra  being  characteristic  for 
each  element.  Moseley  (1914)  showed  that  the  higher  the  atomic  weight  of  the 
material  of  the  anti-cathode  the  shorter  the  wave-length  of  the  characteristic 
X-rays,  and,  further,  for  elements  of  wave-length  between  aluminium  and  gold, 
that  the  wave-length  of  the  stronger  rays  is  simply  connected  with  a  set  of  con- 
secutive integral  numbers  from  13  for  aluminium  to  89  for  gold — numbers,  in  fact, 
which  represent  the  position  of  the  elements  in  the  Periodic  Table,  excluding 
hydrogen.  In  other  words,  within  the  limits  cited,  the  wave-length  of  the  lines 
of  the  X-ray  spectrum  of  each  element  is  connected  in  the  same  way  with  the 
integer  that  represents  the  place  assigned  to  it  by  chemists  in  the  Periodic  Table. 
So  far,  some  elements  of  low  atomic  weight  have  not  been  investigated  by  this 
method,  but  the  rule  is  probably  of  general  application. 

These  results  are  of  particular  interest  in  connection  with  elements  still  missing 
and  those  misplaced  in  the  Periodic  Table.  Thus  the  elements  chlorine  and 
potassium  correspond  with  the  numbers  17  and  19,  leaving  18  for  argon,  not 
determined.  Iron,  nickel  and  cobalt  correspond  with  the  numbers  26,  27,  28. 
It  would  appear  that  only  three  elements  are  missing,  a  rare  earth  element  and 
those  with  the  numbers  43  and  75,  corresponding  with  the  two  vacancies  in  the 
Periodic  Table  below  manganese. 

1  Radio-active  Substances  and  their  Radiations,  1912,  p.  619. 


PROBLEMS   AND   QUESTIONS 

THE  explanations  and  examples  given  in  the  book  should  be  sufficient  to  enable 
the  learner  to  solve  all  the  problems  here  given.  Many  of  the  problems  deal  with 
relationships  between  weight  and  volume  for  gases,  and  a  few  further  remarks  on 
this  question  are  desirable. 

The  first  method  of  expressing  the  relationship  between  weight  and  volume  for 
gases  is  the  following  (p.  no)  : — 

(i)  The  molecular  weight  of  any  gas  in  grams,  at  N.  T.P.,  occiipies  22.4  litres. 

A  second  method  depends  on  the  use  of  hydrogen  as  unit  of  density.  Since  2.016 
grams  of  hydrogen  at  N.T.P.  occupy  22.4  litres,  i  litre  of  hydrogen  weighs 
2.016/22.4=0.09  gram,  and  a  litre  of  any  other  gas  of  density  d  at  N.T.P.  must 
weigh  0.09  x  d  grams.  The  density  of  a  gas  is  obtained  from  its  molecular  weight 
on  the  oxygen  standard  by  dividing  by  2,  or,  more  accurately,  by  2.016  (for  ex- 
ample, the  molecular  weight  m  of  hydrogen  is  2.016,  and  as  the  density  of  the 
gas  is  taken  as  unit  it  is  evident  that  m/2.oi6=d).  Hence  we  have,  as  the  second 
method  of  expressing  the  relationship  between  weight  and  volume  for  gases,  the 
following: — 

(2)  One  litre  of  gas  at  N.  T.P.  weighs  0.09  x  —          '. 

Which  of  these  methods  is  found  most  convenient  in  any  given  case  depends 
upon  the  nature  of  the  problem ;  the  first  is  more  generally  useful.  Neither  of 
them  is  strictly  accurate  (the  error  depends  on  the  deviation  of  the  gas  in  question 
from  the  simple  gas  laws),  and  therefore  calculation  with  approximate  atomic 
weights  is  usually  sufficiently  accurate. 

ELEMENTARY. 

(1)  What  criteria  would  you  use  to  distinguish  between  physical  and  chemical 
changes  ?     Classify  the  following  as  physical  or  chemical  changes  respectively, 
giving  your  reasons :  (a]  the  burning  of  a  candle,  (b)  the  dissolving  of  sugar  in 
water,  (c)  the  magnetizing  of  iron,  (d)  the  rusting  of  iron,  (e)  the  melting  of  ice. 

(2)  If  you  were  supplied  with  a  material  which  might  be  an  element,  a  mechanical 
mixture,  or  a  chemical  compound,  describe  as  clearly  as  possible  how  you  would 
proceed  to  demonstrate  its  nature. 

(3)  In  Lavoisier's  experiment  on  the  oxidation  of  tin  (p.  28),  is  it  to  be  anti- 
cipated that  the  gain  in  weight,  due  to  the  entrance  of  air  when  the  vessel  is 
opened,  will  be  exactly  equal  to  the  gain  in  weight  of  the  tin  during  the  heating? 
Give  reasons  for  your  answer. 

(4)  In  what  manner  was  the  phlogistic  theory  overthrown  by  Lavoisier  ?     Give 
the  arguments  used  in  support  of  that  theory,  and  explain  how  it  was  shown  that 
they  cannot  be  maintained.     (Univ.  Coll.,  London.) 

(5)  If  100  c.c.  of  a  gas,  measured  at  10°  and  740  mm.,  is  warmed  to  100°  and 
the  volume  is  kept  constant,  what  is  the  resulting  pressure? 

(6)  Fifty  c.c.  of  a  gas  at  10°  exerts  a  pressure  of  700  mm.,  to  what  temperature 
must  it  be  heated  in  order  that  it  may  exert  a  pressure  of  1000  mm.,  the  volume 
being  kept  constant  ? 

597 


598     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

(7)  A  certain  quantity  of  a  gas  measures  75  c.c.   at   —20°  C.  and  50x3  mm. 
pressure.     What  will  be  its  volume  (a)  at  — 100°  and  1000  mm.  pressure,  (b)  at 
100°  and  760  mm.  pressure? 

(8)  One  quantity  of  a  gas  measures  100  c.c.  at  10°  and  1000  mm.  pressure  ; 
another  quantity  measures  50  c.c.  at  —20°  and  1600  mm.  pressure.     Which  of  the 
two  quantities  has  the  greater  weight  ? 

(9)  Explain  in  what  manner  the  volume  of  a  gas  varies  (a)  with  change  of  tem- 
perature, (b}  with  change  of  pressure.     Calculate  what  pressure  100  volumes  of  a 
gas  (at  unit  pressure)  will  exert  when  compressed  to  10  volumes,  and  heated  from 
10°  to  100°.     (Univ.  Coll.,  London,  modified.'] 

(10)  An  open  vessel  is  heated  till  one-third  of  the  air  it  contains  at  20°  is  expelled. 
What  is  the  temperature  of  the  vessel  ? 

(n)  If  25  c.c.  of  nitrogen  take  ten  minutes  to  diffuse  from  the  containing  vessel, 
how  long  would  it  take  for  50  c.c.  of  chlorine  to  diffuse  under  the  same  conditions? 

(12)  If  10  c.c.  of  oxygen  diffuse  in  ten  minutes  through  a  crack  in  a  containing 
vessel,  what  volumes  of  ammonia,  nitric  oxide,  and  sulphur  dioxide,  respectively, 
would  pass  out  in  the  same  time  under  similar  conditions  ? 

(13)  Define  the  terms  critical  temperature  and  critical  pressure.     Explain  the 
importance  of  critical  phenomena  in  connection  with  the  liquefaction  of  gases. 
Give  a  brief  account  of  the  general  principles  underlying  the  modern  methods  of 
liquefying  the  so-called  ' '  permanent ' '  gases. 

(14)  State  the  law  of  partial  pressures.    Fifty  c.c.  of  hydrogen  at  10°  and  700  mm. 
and  100  c.c.  of  oxygen  at  30°  and  500  mm.  pressure,  are  mixed  at  15°  in  a  vessel  of 
120  c.c.  capacity.    What  is  the  total  pressure  in  the  vessel  and  the  partial  pressure 
of  each  gas  ? 

(15)  If  a  certain  quantity  of  a  gas  when  dry  exerts  a  pressure  of  700  mm.  at  20°, 
what  will  be  the  total  pressure  in  the  vessel  if  a  little  water  is  added?    (Cf.  p.  65.) 

(16)  Explain  the  difference  between  the  "  absorption  coefficient  "  of  a  gas  and 
its  "  solubility."     The  absorption  coefficient  of  nitrogen  in  water  at  30°  is  0.0138. 
What  is  its  "  solubility  "  at  the  same  temperature? 

(17)  What  are  the  principal  mineral  ingredients  of  natural  waters?    Discuss 
their  importance  from  the  point  of  view  of  (a)  contamination  of  the  water  by 
sewage  or  other  animal  refuse,  (b}  the  use  of  the  water  for  washing  and  boiler 
purposes.     (Birmingham  Univ.) 

(18)  Berzelius  carried  out  the  following  experiments  on  the  three  oxides  of  lead  : 
(a)  10  grams  of  lead  on  heating  gave  10.78  grams  of  litharge,  (b)  9.835  grams  of 
red  lead  after  heating  strongly  gave  9.545  grams  of  litharge,  (c)  4.87  grams  of 
lead  peroxide  lost  on  heating  0.325  gram  of  oxygen,  the  residue  being  litharge. 
Use  these  results  to  test  the  law  of  multiple  proportions. 

(19)  State  the  law  of  multiple  proportions.    An  element  forms  oxides  containing 
22-S3'   27.95,  30.38,  and  36.78  per  cent,  of  oxygen;  without  using  any  atomic 
weights  show  that  this  is  in  accordance  with  the  law,  and  indicate  the  most  probable 
formulae  for  the  oxides,  taking  E  as  the  symbol  of  the  element.      (London  Univ.] 

(20)  Three  binary  compounds  containing  mercury,  an  oxide,  an  iodide,  and  a 
sulphide,  were  found  to  contain  92.59,  61.16,  and  86.18  per  cent,  of  the  metal 
respectively.     What   information   do   these  results  afford  as  to  the   combining 
weight  of  mercury  ? 

(21)  State  as  clearly  as  possible  the  reasons  which  led  Avogadro  to  enunciate 
the  hypothesis  associated  with  his  name.     (Birmingham  Univ.] 

(22)  What  do  you  understand  by  an  empirical,  a  molecular,  and  a  constitutional 
formula?    Illustrate  your  answer  by  an  example  in  each  case.     How  does  the 
formula  of  a  substance  indicate  its  composition  by  weight?     (Conjoint  Board.} 

(23)  Explain  how  it  has  been  shown  that  the  molecular  formula  for  water  is 
H2O,  noting  specially  any  assumptions  made  in  the  course  of  your  proof. 

(24)  Find  the  percentage  amounts  of  the  elements  in  the  following  compounds, 
using  approximate  atomic  weights :    (a)  ferric  chloride,   FeCl3 ;   (b}   potassium 
permanganate,    KMnO4 ;    (c)   cryolite,  AlF3,3NaF ;    (d)   ammonium    sulphate, 
(NH4)2SO4;  (e)  potash  alum,  A12(SO4)3,K2SO4,24H2O. 


PROBLEMS  AND  QUESTIONS  599 

(25)  (a)  What  is  the  weight  of  a  litre  of  each  of  the  gases  chlorine,  hydrogen 
chloride,  nitric  oxide,  NO,  and  carbon  dioxide,  CO2,  at  17°  and  913  mm.  pressure? 
(b)  What  volume  would  be  occupied  by  i  gram  of  each  of  these  gases  at  23°  and 
800  mm.  pressure? 

(26)  One  gram  of  hydrogen  and  one  gram  of  oxygen  are  put  in  a  vessel  of  10 
litres  capacity  at  20°.     What  is  the  total  pressure  of  the  mixture  and  the  partial 
pressure  of  each  of  the  constituents  ? 

If  the  mixture  of  hydrogen  and  oxygen  is  shaken  up  with  water,  in  what  ratio 
would  the  gases  dissolve? 

(27)  The  specific  heat  of  a  metal  is  0.031,  and  3.09  grams  gave  3.33  grams  of 
the  oxide.     Calculate  the  approximate  atomic  weight  from  these  data.     (St. 
Andrews  Univ.) 

(28)  The  two  chlorides  of  a  metal  contain  44.10  and  34.46  per  cent,  of  the  metal 
respectively,  and  its  specific  heat  is  o.uo.     Calculate  the  exact  atomic  weight  of 
the  metal  and  the  formulae  of  the  chlorides. 

(29)  The  vapour  density  of  the  chloride  of  a  metal  m  is  40.96,  and  it  contains 
11.375  per  cent,  of  m.     Find  the  probable  atomic  weight  of  the  element  and  the 
formula  of  the  chloride.     (St.  Andrews  Univ.] 

(30)  0.125  gram  of  a  substance,  when  vaporized  in  a  Victor  Meyer  apparatus,  dis- 
placed a  volume  of  air  which,  when  collected  over  water  at  20°,  measured  26  c.c. , 
the  total  pressure  inside  the  vessel  being  757.5  mm.     Calculate  the  molecular 
weight  of  the  substance. 

(31)  A   piece  of  metal  weighing  i.i  gram,  when  dissolved  in   concentrated 
hydrochloric  acid,  displaces  209  c.c.  of  hydrogen  at  o°  C.  and  760  mm.     The 
specific  heat  of  the  metal  is  0.154.     Calculate  (a)  the  equivalent  weight  of  the 
metal,  (b)  its  atomic  weight.     (Univ.  Coll.,  London.) 

(32)  What  weight  of  zinc  would  be  required  to  furnish  hydrogen  sufficient  to 
fill  a  balloon  capable  of  lifting  its  own  weight  (including  car)  of  50  kilograms  to 
a  height  where  the  atmospheric  pressure  is  500  mm.  and  the  temperature  7°. 
(Density  of  air  14.44  for  density  of  H  =  i.)     (Univ.  Coll.,  London.) 

(33)  Two  gaseous  oxides  of  an  element  contain  respectively  63.7  and  46.7  per 
cent,  of  the  element,  and  :f  litre  of  these  oxides  weighs  1.98  grams  and  1.36  grams 
at  o°  and  760  mm.    Calculate  the  equivalent  and  atomic  weight  of  the  element. 

(34)  1.6944  grams  of  the  chloride  of  a  metal  precipitate  as  chloride  2  grams  of 
silver.     The  specific  heat  of  the  metal  is  0.0567.     Find  (a)  the  equivalent  weight 
of  the  metal,  (b)  its  atomic  weight,  and  (c)  the  formula  for  the  chloride.    (Aberdeen 
Univ.) 

(35)  A  divalent  metal  forms  several  oxides ;  it  was  found  that  when  411  grams 
of  an  oxide  (containing  9.34  per  cent,  of  oxygen)  were  treated  with  nitric  acid  that 
397.2  grams  of  the  nitrate  had  been  formed,  whilst  143.4  grams  of  another  oxide 
(containing  13.39  Per  cent-  °f  oxygen)  were  left  undissolved. 

Calculate  the  formulae  of  the  oxides  and  the  atomic  weight  of  the  metal,  and 
write  out  the  equation  representing  the  reaction  which  took  place.    (London  Univ. ) 

(36)  What  weight  of  common  salt  would  be  required   to   furnish  sufficient 
hydrogen  chloride  to  neutralize  100  grams  of  a  30  per  cent,  solution  of  caustic 
soda?    What  would  be  the  volume  of  the  gas  at  N.T.P.  ?    (Board  of  Education.) 

(37)  When  excess  of  silver  nitrate  was  added  to  a  solution  of  sodium  chloride 
2.87    grams    of  silver   chloride   were   precipitated.      Calculate    the   amount   of 
(a)  sodium  chloride,  (b)  combined  chlorine,  in  the  original  solution.     [Equation: 
NaCl  +  AgNO3= AgCl  +  NaNO3.] 

(38)  State  concisely  what  information  is  afforded  by  the  following  symbols  and 
equations :-  2H2+O2=2H2O. 

AgNO3,Aq  +  NaCl,Aq  =  AgCl  +  NaNO3,Aq.     (Birmingham  Univ.) 

(39)  A  compound  of  carbon  and  hydrogen  contains  85.7  per  cent,  of  the  former 
element  and  its  vapour  density  is  14.    Find  its  empirical  and  molecular  formulae. 

(40)  A  salt  contains  33.86  per  cent,  of  copper,  14.93  per  cent,  of  nitrogen,  and 
51.20  per  cent,  of  oxygen.     What  is  its  empirical  formula? 


6oo     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

(41)  A  substance  contains  32.79  per  cent,  sodium,  13.02  per  cent,  aluminium, 
and  54.19  per  cent,  fluorine.     Find  its  simplest  formula,  using  accurate  atomic 
weights.     (  Board  of  Education.  ) 

(42)  A  crystalline  solid  contains  5.7  per  cent,  of  aluminium,  8.2  per  cent,  of 
potassium,  13.5  per  cent,  of  sulphur,  5.1  per  cent,  of  hydrogen,  the  remainder 
being  oxygen.     Express  its  composition  in  terms  of  the  groups  Al,.O3,  K2O,  SO*, 
and  H2O. 

(43)  On  mixing  solutions  of  potassium  sulphate  and  nickel  sulphate,  crystals 
separate  which  contain  both  metals  and  on  analysis  are  found   to   have  the 
composition:  K  =  i7.8,  Ni  =  i3.5,  SO4=44,  and  H2O—  24.7  per  cent.     What  is 
the  formula  of  the  salt  ?    (Birmingham  Univ.  ) 

[One  gram  of  a  mixture  of  potassium  andi  sodium  chlorides  on  treatment  with 
excess  of  silver  nitrate  gave  2.1  grams  of  silver  chloride.  What  was  the  proportion 
of  the  two  salts  in  the  original  mixture  ? 

As  problems  of  this  type  are  not  considered  in  the  book,  it  is  desirable  to  give 
a  solution  of  the  above  example.  The  equations  concerned  are  :  — 

NaCl  -»  AgCl  and  KCl-»AgCl, 
58.5-»i43-5          74.5^143.5. 

the  reacting  quantities  being  calculated  as  usual  from  the  atomic  weights. 
If  x  is  the  amount  of  KC1  in  the  mixture,  then  x  grams  KC1  give 

x  x  I43'5  grams  AgCl. 

74-5 
Similarly,  since  i  —  x  is  the  amount  of  NaCl  present,  i  —  x  grams  NaCl  give 

i  -  x  x  I43'5  grams  AgCl. 

58-5 
The  total  amount  of  silver  chloride  is  2.1  grams,  therefore 


74-5  58.5 

Solving,  we  obtain  ^=0.67,  hence  the  mixture  contains  67  per  cent.  KC1  and 
33  per  cent.  NaCl.] 

As  a  further  illustration  of  this  method  of  indirect  analysis  the  following  problem 
should  be  worked  out  :  — 

(44)  1.5  grams  of  a  mixture  of  sodium  and  potassium  chlorides  were  converted 
to  the  normal  sulphates  by   evaporation   with   sulphuric   acid,  the  mixture   of 
sulphates  weighing  1.798  grams.     What  was  the  composition  of  the  mixture  ? 

(45)  Mention  some  processes,  other  than  the  action  of  the  electric  discharge  on 
oxygen,  by  which  ozone  is  produced.     On  partly  ozonising  100  c.c.  of  oxygen  a 
decrease  of  volume  of  5  c.c.  was  observed.     What  volume  of  gas  would  be  left, 
and  what  weight  of  iodine  liberated,  on  treating  the  ozonised  oxygen  with  excess 
of  potassium  iodide? 

(46)  What  do  you  understand   by   the  heat  of  formation   of  a   compound? 
Indicate  the  methods  available  for  the  determination  of  the  heat  of  formation  of 
compounds.     (Aberdeen  Univ.] 

(47)  Calculate  the  heat  of  formation  of  ethylene  from  its  elements  at  constant 
pressure  from  the  following  data:  Heat  of  combustion,  333,350  cal.  ;  heats  of 
formation  of  carbon   dioxide,   94,300   cal.,    and   of  liquid   water,   68,400   cal. 
[Equation:  C2H4+3O2=2CO2+2H2O.] 

(48)  Calculate  the  heat  evolved  in  the  combination  of  ethylene  and  hydrogen 
to  ethane  from  the  following  data:  Heat  of  combustion  of  ethane,  370,440  cal.  ; 
of  ethylene,  333,350  cal.  ;  of  hydrogen,  68,400  cal. 

(49)  By  what  reactions  is  chlorine  usually  prepared  in  the  laboratory?     State 
the  chief  properties  of  the  element,  and  point  out  in  what  respects  it  resembles  and 
differs  from  the  other  halogens.     (Univ.  Coll.,  London.} 

(50)  Compare  the  halogen  acids  in  respect  to  modes  of  formation  and  stability. 
(St.  Andrews  Univ.} 


PROBLEMS  AND  QUESTIONS  601 

(51)  State  the  theorem  of  Le  Chatelier  and  apply  it  to  the  discussion  of  the 
influence  of  pressure  on  the  equilibrium  N2O4<^2NO2.     (Aberdeen  Univ.) 

(52)  When  steam  is  passed  over  red-hot  iron,  hydrogen  and  oxide  of  iron  are 
formed  ;  and  when  hydrogen  is  passed  over  heated  oxide  of  iron,  steam  and  metallic 
iron  are  produced.     Explain  the  apparent  discrepancy,  and  give  a  brief  account 
of  the  general  principles  thereby  illustrated.     (Aberdeen  Univ.,  modified.) 

(53)  Describe  three  processes  for  the  approximate  determination  of  molecular 
weights   of  substances  in  aqueous  solution.     Give  a  sketch   of  the  apparatus 
required  in  each  instance,  and  explain  how  the  results  obtained  are  to  be  inter- 
preted.    (Birmingham  Univ.) 

(54)  1.2  grams  of  a  substance  dissolved  in  24.5  grams  of  water  (K  =  i8.s) 
caused  a  depression  of  the  freezing-point  of  1.05°.     Find  the  molecular  weight  of 
the  substance. 

(55)  The  addition  of  1.065  grams  of  iodine  to  30.14  grams  of  ether  (K=2i.o) 
raises  the  boiling-point  of  the  latter  by  0.296°.     What  is  the  molecular  weight  of 
iodine  in  ether? 

(56)  What  would  be  the  osmotic  pressures  in  aqueous  solutions  of  (a)  hydrogen, 

(b)  nitrogen,  (c)  carbon  dioxide,  saturated  at  o°  and  76  cm.  pressure? 

(57)  What  are  the  principal  atmospheric  gases,  and  what  part  do  they  play  in 
the  life  of  animals  and  plants.     (Univ.  Coll.,  London.) 

(58)  An  analysis  of  impure  air  was  made  by  absorbing  the  carbon  dioxide  with 
potash  and  the  oxygen  with  phosphorus.     The  volume  was  kept  constant,  but 
pressure  and  temperature  varied  during  the  reading  of  the  volume. 

Original  volume,  10  c.c.  at  17°  C.  and  740  mm. 
After  treatment  with  caustic  potash,  10  c.c.  at  18°  and  720  mm. 
After  removal  of  oxygen  with  phosphorus,  19  c.c.  at  17°  and  400  mm. 
State  the  result  in  volumes  per  cent.     (Univ.  Coll.,  London.) 

(59)  What   happens   in  each  case  when  air  containing  one   of  the  following 
gases  is  passed  over  starch-potassium  iodide  papers:    (a)  chlorine,  (b) ozone,  (<•) 
nitrogen  peroxide?     Explain  the  reactions  in  each  case.     (St.  Andrews  Univ.) 

(60)  What  are  the  chief  sources  of  ammonia  at  the  present  time,  and  for  what 
purposes  is  ammonia  chiefly  used?     What  explanation   can  be   given  of  the 
development  of  heat  when  ammonia  is  dissolved  in  water ;  what  is  the  compo- 
sition of  the  aqueous  solution,  and  what  is  the  effect   of  boiling  the  solution? 
(London  Univ.) 

(61)  How  can  it  be  shown  that  the  vapour  obtained  by  heating  ammonium  chloride 
consists  of  a  mixture  of  ammonia  and  hydrogen  chloride  gases?    What   other 
decompositions  of  a  similar  nature  are  you  acquainted  with  ?    (Birmingham  Univ.) 

(62)  When  60  c.c.  of  a  mixture  of  nitrous  and  nitric  oxides  are  mixed  with 
excess  of  hydrogen  and  the  mixture  exploded,  40  c.c.  of  nitrogen  remain.     What 
was  the  composition  of  the  mixture  ? 

(63)  Write  a  short  essay  on  oxidation  and  reduction,  giving  appropriate  illustra- 
tions in  explanation.     (St.  Andrews  Univ.] 

(64)  You  are  supplied  with  a  source  of  nitric  oxide  in  large  quantity.    Carefully 
describe  how  you  would  prepare  from  it  (a)  concentrated  nitric  acid,  (b)  ammonia, 

(c)  hydrazoic  acid,  (d)  liquid  nitrogen  peroxide.     (Birmingham  Univ.) 

(65)  Give  an  account  of  the  various  means  by  which  atmospheric  nitrogen  can 
be  brought  into  combination  for  industrial  purposes.     (London  Univ.) 

(66)  What  different  compounds  are  obtainable  by  the  action  of  metals  on 
nitric  acid.      Give  equations  and  indicate  the  necessary  conditions  in  each  case. 
(Birmingham  Univ.) 

(67)  Explain  briefly  what  is  meant  by  the  valency  of  an  element.     Is  the  valency 
of  an  element  always  the  same?     If  not,  how  does  it  usually  vary?     Discuss  the 
valency  of  carbon,  nitrogen,  oxygen,  chlorine,  and  iron.    (London  Univ.) 

(68)  What  do   you   understand   by  the  term   "reversible  reaction"?      Give 
examples  and  show  how  such  reactions  may  be  made  practically  complete  in 
either  direction.     (King's  College,  London.) 

(69)  Write  a  short  account  of  the  phenomenon  generally  known  as  "  catalysis." 


602     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

Mention  substances  which  act  catalytically  and  state  two  cases  which  are  of 
commercial  importance.     (King's  College,  London.} 

(70)  32  c.c.  of  a  solution  of  sodium  hydroxide  were  found  to  neutralize  28  c.c. 
of  0.8   N   hydrochloric  acid.     Calculate  the  volume  of  water  which  must  be 
added  to  a  litre  of  the  sodium  hydroxide  solution  in  order  to  make  it  exactly 
0.5  N.     (St.  Andrews  Univ.] 

(71)  Discuss  the  mechanism  of  the  process  of  electrolysis.     Illustrate  your 
answer  by  reference  to  the  electrolysis  of  dilute  sulphuric  acid,  of  fused  sodium 
hydroxide,  and  of  sodium  chloride  solution.     (St.  Andrews  Univ.] 

(72)  What  do  you  understand  by  a  decinormal  solution  ?    Calculate  the  weight 
of  the  solute   per  litre  in  decinormal  solutions  of  the   following:    (a)  sodium 
carbonate,   (6)   orthophosphoric  acid,  (c)  barium  hydroxide,  (d)  silver  nitrate, 
(e)  iodine,  (f)  potassium  dichromate. 

(73)  Define  the  terms  (a)  basic  salt,  (b)  weak  acid,  (c)  strong  base,  (d)  ionisa- 
tion,  (e)  hydrolysis.     Give  examples  of  each.     (Birmingham  Univ.] 

(74)  What  do  you  understand  by  the  term  ion  ?     Illustrate  the  part  played  by 
the  ions  in  electrolysis  by  reference  to  (a)  the  electrolytic  decomposition  of  hydro- 
chloric acid,  (l>)  the  electrodeposition  of  copper.     (6V.  Andrews  Univ. ,  modified. ) 

(75)  What  are  basic  salts?     Illustrate  your  answer  by  reference  to  at  least  three 
definite  examples,  and  explain  how  you  would   ascertain   whether  a  definite 
substance  was  a  ncrmal,  acid,  or  basic  salt. 

(76)  How  many  types  of  oxide  exist  ?     Give  two  examples  of  each  type,  and 
show  how  they  differ  in  chemical  behaviour,  as,  for  example,  towards  (a)  water, 
(b}  hydrochloric  acid,  (c)  caustic  soda.     (Sheffield  Univ.) 

To  what  classes  would  you  assign  the  following  oxides  :  (a)  CO,  (b)  Pb3O4, 
(c)  CrO3,  (d)  NagOa.     (Birmingham  Univ.) 

(77)  Mention  three  general  methods  by  which  metals  may  be  converted  into 
their  oxides.     Illustrate  each  method  by  describing  an   experiment.     (London 
Univ. ) 

(78)  What  weight  of  sodium  hydroxide  would  be  required  to  neutralize  TOO  c.c. 
of  a  normal  solution  of  sulphuric  acid  ?    What  weight  and  what  volume  of  carbon 
dioxide,  measured  at  15°  and  760  mm.,  would  be  required  to  convert  completely  the 
above  amount  of  sodium  hydroxide  into  sodium  bicarbonate.     (Conjoint  Board.) 

(79)  A  quantity  of  ammonium  chloride  is  boiled  in  an  open  vessel  with  100  c.c.  of 
normal  sodium  hydroxide  until  no  further  chemical  change  takes  place  ;  it  is  then 
found  that  the  excess  of  sodium  hydroxide  requires  10  c.c.  of  normal  sulphuric 
acid  to  neutralize  it.     How   much  ammonium  chloride  was  used?    (Conjoint 
Board. ) 

(80)  By  means  of  what  reactions  and  tests  can  the  following   gases  be  dis- 
tinguished from  one  another  :  (a)  chlorine  from  sulphur  dioxide,  (b)  nitrogen  from 
carbon  dioxide,  (c)  oxygen  from  nitrous  oxide  ?     (Conjoint  Board.) 

(81)  What  do  you  understand  by  the  electrochemical  equivalent  of  an  element  ? 
If  the  same  quantity  of  electricity  is  passed  through  the  following  solutions : 
(a)  caustic  potash,   (b)  silver  nitrate,  (c)   cuprous  chloride,    (d)  auric  chloride, 
(e)  zinc  chloride,  state  what  substances  are  liberated  at  the  cathode  and  the  relative 
proportion  by  weight  in  which  they  are  formed. 

(82)  Describe  what  occurs  when  electric  sparks  are  passed  through  (a)  hydrogen 
sulphide,  (b)  ammonia,  (c)  air,  noting  also  what  changes  in  volume,  if  any,  are 
observed.     (Birmingham  Univ.) 

(83)  How  would  you  obtain  the  following  elements  from  their  respective  com- 
pounds:   (a)  carbon  from  carbon  dioxide,   (b)  sulphur  from  hydrogen  sulphide, 
(c)  hydrogen  from  ammonia,  (d)  chlorine  from  hydrogen  chloride,  (e)  nitrogen 
from  nitric  oxide. 

(84)  What  experiments  would  you  make  in  order  to  show  that  the  following 
substances  are  chemical  compounds  and  not  simple  substances  or  mixtures:  (a) 
water,  (b)  carbon  dioxide,  (c)  nitrous  oxide,  (d)  ferrous  sulphide,  (e)  cupric  oxide. 
(Conjoint  Board.) 

(85)  Explain  what  is  meant  by  combustion,  both  "  complete  "  and  "  incomplete.' 


PROBLEMS  AND  QUESTIONS  603 

Compare  and  contrast  the  flames  of  a  candle,  a  blowpipe,  and  a  Bunsen  burner 
(both  with  the  air  admitted  and  with  the  air  shut  off).     {London  Univ.) 

(86)  What  volume  of  oxygen,  at  N.T.P.,  will  be  required  for  the  complete  com- 
bustion of  i  gram  of  carbon  monoxide,  and  what  will  be  the  volume  of  the 
resulting  carbon  dioxide.,  measured  at  20°  and  760  mm.  pressure? 

(87)  What  would  be  the  volume,  at  13°  and  740  mm.  pressure,  of  each  of  the 
products  formed  by  the  complete  combustion  of  114  grams  of  carbon  disulphide? 
What  amount  of  caustic  soda  would  be  saturated  by  these  products  if  (a)  the 
normal  salts,  (b)  the  acid  salts,  were  formed?     (London  Univ.) 

(88)  What  would  be  the  volume  and  composition  of  the  gases  resulting  from 
the  explosion  of  a  mixture  of  10  c.c.  of  hydrogen  and  10  c.c.  of  methane  with 
35  c.c.  of  oxygen? 

(89)  A  sample  of  gas  had  the  following  percentage  composition:  hydrogen  45, 
marsh  gas  30,  carbon  monoxide  20,  acetylene  5.     100  volumes  of  it  were  mixed 
with  160  volumes  of  oxygen,  and  the  mixture  exploded.     Calculate  the  volume 
and  the  composition  of  the  resulting  mixture  of  gases  (all  being  supposed  dry). 
(London  Univ.) 

(90)  Describe  fully  how  the  heat  of  combustion  of  carbon  can  be  accurately 
determined. 

The  combustion  of  a  given  weight  of  carbon  in  oxygen  gives  rise  to  the  same 
amount  of  heat  as  when  it  is  burnt  in  air,  but  considerably  more  heat  is  given 
out  when  it  is  burnt  in  nitrous  oxide.  Explain  these  facts.  (Birmingham 
Univ. ) 

(91)  What  do  you  understand  by  the  terms  "acidic"  and  "basic"  oxides? 
State  to  which  class  each  of  the  following  oxides  belongs  and  give  your  reasons  : 
(a)  magnesium  oxide,   (£)   boron  trioxide,   (c)  alumina,   (d)  sulphur  dioxide,  (e) 
chlorine  monoxide.     (Conjoint  Board.) 

(92)  In  what  ways  are  acids  prepared  from  their  salts?     Illustrate  your  answer 
by  reference  to  hydrochloric  acid,  phosphoric  acid,  and  boric  acid.    {Birmingham 
Univ.) 

(93)  Starting  with  ordinary  quartz  sand,  how  would  you  prepare  (a)  soluble 
silicic  acid,  (b)  insoluble  silicic  acid,  (c)  silicon  dioxide?     (Aberdeen  Univ.) 

(94)  Devise  a  series  of  reactions  by  means  of  which  specimens  of  the  con- 
stituent elements  could  be  obtained  from  potassium  chlorate.      (St.   Andrews 
Univ.) 

(95)  How  would  you  detect  the  following  impurities,  not  more  than  i  per  cent, 
being  present  in  each  case:  (a)  oxygen  in  a  sample  of  nitrogen,  (b)  ammonia  in 
water,  (c)  potassium  carbonate  in  caustic  potash,  (d)  potassium  chloride  in  nitre, 
(e)  chlorine  in  hydrochloric  acid,  (f)  sulphur  dioxide  in  carbon  dioxide?    (London 
Univ. ) 

(96)  Describe  and  explain  the  phenomena  of  efflorescence  and  deliquescence. 
(Aberdeen  Univ.) 

(97)  Explain  the  following  in  terms  of  the  theory  of  electrolytic  dissociation: 
(a)  the  neutralization  of  acids  by  alkalis,  (b)  double  decomposition,  (c)  the  colour 
of  solutions  of  copper  salts.     (St.  Andrews  Univ. ) 

(98)  What   are   alloys   and    how  are   they   prepared?      Give  an  account  of 
analytical  tests  which  you  would  employ  to  identify  the  metais  present  in  a  silver 
coin.     (St.  Andrews  Univ.) 

(99)  Given  ordinary  standard  silver  (which  contains  7^  per  cent,   of  copper), 
how  would  you  prepare  pure  silver  nitrate?     How  does  silver  nitrate  react  with 
solutions  of  (a)  hydrogen  sulphide,  (b)  caustic  soda,  (c)  common  sodium  phosphate, 
(d)  potassium  iodide,  (e)  hydrocyanic  acid?    State  any  change  in  the  colour  of  the 
solution  both  before  and  after  the  addition  of  nitric  acid.     (London  Univ.) 

(100)  Discuss  the  following  statements:  (a)  oxidation  is  equivalent  to  an  in- 
crease of  the  positive  or  a  decrease  of  the  negative  ionic  charge ;  (b)  the  valency 
of  an  element  is  numerically  equal  to  the  number  of  univalent  atoms  or  groups 
which  can  combine  with  one  atom  of  the  element.     {Birmingham  Univ.) 

(roi)  State  the  chief  uses  of  lime  and  explain  the  part  it  plays  in  (a)  the  soften- 


604     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 

ing  of  hard  water,  (l>)  the  hardening  of  mortar,  (c)  the  formation  of  blast-furnace 
slag.     (Univ.  Coll.,  London.) 

(102)  Starting  from  dolomite,  describe  the  preparation  of  pure  specimens  of 
magnesium  sulphate,  magnesium  oxide,  and  metallic  magnesium.     (St.  Andrews 
Univ. ) 

(103)  0.5  gram  of  a  mixture  of  calcium  and  magnesium  carbonates  lost  on 
heating  o.  24  gram.     Calculate  the  proportion  of  each  carbonate  in  the  mixture. 
(Birmingham  Univ. ) 

(104)  Discuss  the  following  changes :  (a)  the  precipitation  of  barium  chloride 
from  solution  by  concentrated  hydrochloric  acid ;  (b)  the  action  of  concentrated 
sulphuric  acid  on  mercuric  chloride  ;  (c)  the  efflorescence  of  washing-soda  ;  (d)  the 
solubility  of  magnesium  hydroxide  in  ammonium  chloride  solution.      (London 
Univ.) 

(105)  Compare  and  contrast  the  properties  of  the  most  important  compounds 
of  calcium  and  magnesium.     (St.  Andrews  Univ.) 

(106)  What  is  meant  by  the  statement  that  the  carbonates  of  calcium  and 
magnesium  are  isomorphous?     Give  any  other  instances  of  isomorphism  with 
which  you  are  acquainted.     (Birmingham  Univ. ) 

(107)  Give  the  names  and  formulae  of  the  oxides  of  carbon,  silicon,  tin,  and 
lead,   and  from  a  consideration  of  the  properties  of  these  oxides,  or  some  of 
them,  demonstrate  the  probability  that  the  elements  of  which  they  are  the  oxides 
belong  to  the  same  group  or  family.       (Sheffield  Univ. ) 

(108)  Starting  from  silica,  how  would  you  prepare  (a)  silicon  hydride,  (b)  silicon 
tetrachloride,  (c)  potassium  silicofluoride,  (d)  silicon  chloroform,  (e)  orthosilicic 
acid.      Give  a  brief  description  of  these   substances  and   compare  them  with 
analogous  compounds  of  elements  in   the  same  group   of  the  periodic  table. 
(Dublin  Univ.) 

(109)  Give  an  instance  in  which  each  of  the  following  substances  acts  either  as 
an  oxidizing  or  a  reducing  agent:  (a)  nitric  acid,  (b)  sulphurous  acid,  (c)  strong 
sulphuric  acid,  (d)  hydrogen  peroxide,  (e)  stannous  chloride,  (/)  silver  nitrate. 
Give  equations.     (Conjoint  Board.) 

(no)  How  is  metallic  tin  acted  upon  by  a  solution  of  hydrochloric  acid?  How 
would  you  prove  that  both  the  products  of  the  reaction  are  reducing  agents? 
By  what  processes  could  you  obtain  them  in  an  anhydrous  condition?  (London 
Univ.) 

(in)  What  is  the  nature  of  the  experimental  evidence  in  support  of  the  view 
that  the  molecules  of  argon,  chlorine,  and  arsenic  are  respectively  monatomic, 
diatomic,  and  tetratomic.  (King's  Coll.,  London.) 

(112)  Describe  what  would  be  observed  on   passing  a  current   of  hydrogen 
sulphide  through  solutions  containing  (a)  copper  sulphate,  (b)  arsenious  acid,  (c) 
ferric  chloride,  (d)  zinc  chloride,  each  solution  being  acidified  with  hydrochloric 
acid.     Give  equations  expressing  the  reactions  taking  place.     (Conjoint  Board.) 

(113)  Explain  the  difference  between  wrought  iron,  cast  iron,  and  steel.     Give 
a  brief  account  of  one  method  of  preparing  steel  from  (a)  cast  iron,  (b)  wrought 
iron.     (St.  Andrews  Univ.) 

(114)  What  chemical  industries  may  be  developed  in  a  country  rich  in  natural 
water-power,  where  the  available  minerals  are  coal,  limestone,  and  iron  pyrites? 
(St.  Andrews  Univ.) 

ADVANCED. 
[The  questions  are  not  arranged  in  any  special  order.] 

(1)  Describe  two  methods,  one  of  them  chemical  and  the  other  physical,  for  the 
separation  of  oxygen  from  the  air.     How  could  the  purity  of  the  oxygen  thus 
obtained  be  determined?     (Sheffield  Univ.) 

(2)  State  in  full  all  the  reasons  which  occur  to  you  why  the  formulstof  hydrogen 
sulphide  is  written  as  H2S.     (Univ.  Coll.,  London.) 

(3)  Give  a  sketch  of  the  relationship  between  the  hydrides  of  nitrogen,  showing 


PROBLEMS  AND  QUESTIONS  605 

how  each   may  be  prepared,  and   describing  their   properties.*  (Univ.   Coll., 
London.} 

(4)  What  is  the  constitutional  formula  for  sulphuric  acid  ?    Why  is  this  formula 
adopted?     Give  the  constitutional  formulae  of  potassium  dichromate  and  of 
chromyl  chloride,  also  of  basic  lead  chromate.     Write  an  equation,  using  con- 
stitutional formulas,  showing  the  action  of  sulphuric  acid   on  the  last-named 
compound.     (Univ.  Coll.,  London.} 

(5)  What  is  the  chief  source  of  ammonia,  and  how  is  the  pure  ammonia  solution 
of  commerce  obtained  from  it?     Why  is  this  solution  said  to  be  one  of  ammonium 
hydroxide,  and  why  is  this  base  said  to  be  weak  in  comparison  with  potassium  or 
sodium  hydroxide?     (Sheffield  Univ.} 

(6)  Describe  and  explain  the  reactions  involved  in  separating  (a)  iron  and 
aluminium,  (b}  mercury  and  lead,  in  qualitative  analysis.     How  can  each  of  these 
metallic  radicals  be  identified  by  a  single  test  in  the  dry  way?     (Sheffield  Univ:) 

(7)  Discuss  the  properties  of  peroxides  and  of  peracids  with  reference  to  their 
bearing  on  the  questions  of  constitution  and  valency.     (St.  Andrews  Univ.) 

(8)  How  are  the  percarbonates ,  perchlorates,  permanganates,  and  persulphates 
obtained?    Give  what  you  consider  to  be  the  constitutional  formulas  of  these  sub- 
stances, and  add  any  remarks  on  the  subject  of  valency  bearing  on  this  question. 
(Aberdeen  Univ.} 

(9)  Give  some  account  of  the  reasons  which  led  to  the  adoption  of  the  usual 
modern    atomic  weight  in   preference   to   the  equivalents   in   use  before   1860. 
(Univ.  Coll.,  London.} 

(10)  What  are  the  exceptions  to  Dulong  and  Petit's  law?     Give  an  outline  of 
the  work  which  has  been  done  to  find  how  far  the  law  can  be  regarded  as  valid. 
(Sheffield  Univ. ) 

(n)  What  is  meant  by  the  term  molecular  heat  ?  [Sp.  heatxmol.  wt.]  Calcu- 
late the  specific  heat  of  solid  oxygen  from  the  following  data  :  sp.  heat  of  alumina, 
0.2195  ;  sp.  heat  of  aluminium,  0.2189  ;  Alt=27.i.  (Sheffield  Univ.] 

(12)  How  have  air,  hydrogen,  and  ozone  respectively  been  obtained  in  the 
liquid  state  ?    Give  a  clear  account  of  the  principles  involved  and  of  the  apparatus 
employed.     (Birmingham  Univ.} 

(13)  Discuss  the  different  classes  of  double  salts  from  the  point  of  view  of  their 
analytical  reactions.     Mention  cases  where  these  differences  are  made  use  of  in 
effecting  the  separation  of  similar  elements.     (Glasgow  Univ.} 

(14)  Write  a  short  description  of  the  periodic  classification  with  special  refer- 
ence to  the   so-called  long  and  short   periods,   valency    and    atomic  volume. 
(Glasgow  Univ.} 

(15)  What  is  the  law  of  mass  action?     How  does  it  apply  (a)  to  systems  con- 
taining solids  and  (b}  to  reversible  reactions  ?    Discuss  the  conditions  of  equilibrium 
in  the  dissociation  of  (i.)  hydrogen  iodide,  (ii.)  nitrogen  peroxide,  (iii.)  a  solution 
of  common  salt.     (Glasgow  Univ.} 

(16)  Describe  the  principal  methods  of  extracting  silver  from  argentiferous  lead. 
State  the  principles  on  which  these  methods  are  based.    (Glasgow  Univ.} 

(17)  Briefly  discuss  all  the  methods  you  are  acquainted  with  for  causing  the 
element  nitrogen  to  enter  into  combination.     (Glasgow  Univ.) 

(18)  Name  four  salts  containing  the  same  elements,  but  in  different  proportions, 
which  can  be  obtained  directly  from  sodium  sulphite?    Explain  how  the  pre- 
paration of  each  of  them  can  be  effected,  and  state  the  evidence  on  which  the 
constitutional  formulas  of  any  two  of  them  are  based.     (London  Univ.) 

(19)  Explain  the  reason  why  chlorine  was  regarded  as  an  oxide.     Summarise 
the  experimental  evidence  which  led  chemists  to  abandon  this  view.     (Birming- 
ham Univ.) 


606     A  TEXT-BOOK  OF  INORGANIC  CHEMISTRY 


ANSWERS 


(5)  975- 3  mm. 

(6)  131.3°  C. 

(7)  (a)  25.6  c.c.  ;  (f>)  72.7  c.c. 

(8)  The  first. 

(9)  13.2  times  the  original  pressure, 
jioj  117.6°  C. 

(11)  32  minutes. 

(12)  13.8    c.c.    NH3,    10.3    c.c.    NO, 

7.1  c.c.  SO2. 

(14)  Total  pressure,  692.8  mm.     Pres- 

sure of  H2,  296.8  mm.  ;  of  O2, 
396  mm. 

(15)  7*7-4  mm- 

(16)  0.0153. 

(20)  Combining  weights  100  and  200. 

(24)(«)Fe,  34.46;  Cl,  65.55.  (*)  K, 
24.7  ;  Mn,  34.8  ;  O,  40.5.  (c)  Na, 
32.86  ;  Al,  12.85  ;  F,  54.29.  (d) 
N,  21.2;  H,  6.1;  S,  24.2;  O, 
48.4.  (e)  K,  8.2;  Al,  5.7;  S, 
13.5;  H,  5.1;  O,  67.5. 

(25)  (a)  C12,  3.55  grams;  HC1,  1.825 

grams;    NO,   1.5  grams;    CO2, 

2.2  grams,      (b)  C12,  329  c.c.  ; 
HC1,  641   c.c.  ;    NO,  779  c.c.  ; 
CO2,  531  c.c. 

(26)  Total  pressure,  963  mm. ;  H2,  906 

mm.  ;  O2,  57  mm.  Ratio  of 
volumes  dissolved :  hydrogen  : 
oxygen =6. 6  :  i. 

(27)  206. 

(28)  55.9  ;  MC12  and  MC13. 

(29)  9.1 ;  MCLj. 

(30)  119. 

(32J  121,600  grams. 
(33)  Equivalents   14  and  7.     At.   wt. 
14. 


(34)  (a)  56;  (6)  112;  (c)  MC12. 

(35)  M3O4  and   MO2  ;    207  ;    M3O4  + 

4HN03=2M(N03)2+M02. 

(36)  43- 875  grams;  16. 8  litres. 

(37)  (a)  1. 17  grams;  (b)  0.71  gram. 
39    CH2;  C2H4. 

(40    Cu(N03)2. 

(41    Na3AlF6. 

(42    K20,  A1203,  4S03,  24H20. 

(43)  NiSO4,  K2SO4l  6H2O. 

(44)  i  gram  NaCl,  0.5  gram  KCL 

(45)  85  c.c.  ;  0.0567  gram. 

(47)  -795ocal. 

(48)  3i,3iocal. 

(54)  86. 

(55)  251. 

(56)  H2,   0.0203    atmos.,    N2,   0.0239; 

CO2,  1.7134  atmos. 
(58)  CO2)  3  vols.  ;  O2)  43  vols. 
(62)  40  c.c.  NO  ;  20  c.c.  N2O. 
(70)  400  c.c. 
(72)  Na2CO3,  5.3  grams;  H3PO4,  3.27 

grams;    Ba(OH)2,  8.55  grams; 

AgNO3,    17    grams ;     I2,     12.7 

grains  ;   K2Cr2O7,  4.92  grams. 

(78)  4  grams  NaOH  ;  2.36  litres  COa. 

(79)  4.815  grams. 

(81)  H,    i.oob;    Ag,    108;    Cu,  63.4; 
Au,  65.7 ;  Zn,  32.5. 

(86)  0.4  litre  O2  ;  0.86  litre  COa. 

(87)  36.15  litres  CO2 ;  72.3  litres  SO2 ; 

normal  salts"  360    grams,    acid 
salts  180  grams. 

(88)  10  c.c.  CO2,  25  c.c.  steam,  and  10 

c.c.  oxygen. 

(89)  Total  volume,    225;     H2O,    no; 

CO2l  60 ;  O2,  55. 
(103)  0.25  gram  of  each. 


(n)  (Advanced)  0.218. 


INDEX 


ABSOLUTE  temperature,  43. 

—  zero,  43,  49. 
Absorption  coefficient,  79. 

—  of  gases  by  charcoal,  329. 
Accumulator,  lead,  507. 
Acetic  acid,  350. 
Acetylene,  348. 

Acid  anhydrides,  176,  186. 
Acidic  oxides,  99,  186,  254,  273. 
Acids,  activity  of,  187,  266. 

—  basicity  of,  187,  254,  276. 

—  general  properties" of,  98,  187,  265. 

—  hydrogen  theory  of,  275. 

—  oxygen  theory  of,  20,  275. 

—  strength  of,  187,  263,  265. 

—  and  salts,  nomenclature,  188. 
Actinium,  592. 
Actinometer,  93. 

Active  mass,  167. 
Adsorption,  329,  482. 
Affinity,  chemical,  15. 

—  free,  272,  579. 
Agate,  368. 

Air,  combustion  in,  27. 

—  liquid,  72,  209. 
Alabaster,  452. 
Alcohol,  definition,  346. 

—  ethyl,  349. 
Aldehydes,  350. 

Alkali  metals,  comparison  of,  422. 

—  waste,  291,  403. 
Allotropic  modifications,  175. 
Alloys,  397,  427,  481. 
Alum,  burnt,  485. 

—  chrome,  535. 
Aluminates,  483. 
Aluminium,  479. 

—  alloys,  481. 

—  bronze,  427,  481. 

—  chloride,  484. 

—  hydroxide,  482. 

—  oxide,  482. 

—  silicates,  370,  485. 
' —  sulphate,  484. 

—  sulphide,  486. 
Alums,  484. 


607% 


Alums,  pseudo,  485. 
Alunite,  485. 
Amalgams,  472. 
Amethyst,  482. 
Amides,  316,  343. 
Ammonia,  213. 

—  composition  of,  217. 

—  soda  process,  405. 
Ammoniacal  liquor,  214,  418. 
Ammonium,  421. 

—  amalgam,  421. 

—  carbamate,  420. 

—  carbonates,  420. 

—  chloride,  419. 

dissociation  of,  215,  419. 

—  dichromate,  202,  538. 

—  hydroxide,  216,  419. 

ionic  dissociation  of,  263,  419. 

—  magnesium  arsenate,  518. 
phosphate,  464. 

—  molybdate,  540. 

—  nitrate,  233. 

—  nitrite,  201. 

—  phosphomolybdate,  540. 

—  salts,  418. 

—  stannichloride,  499. 

—  sulphate,  420. 

—  sulphides,  421. 

—  thioantimonate,  527. 

—  thioantimonite,  526. 

—  thioarsenate,  519. 

—  thioarsenite,  519. 

—  thiostannate,  499. 
Analysis,  gravimetric,  287. 

—  volumetric,  285. 
Anatase,  492. 
Andalusite,  370. 
Anglesite,  500. 
Anhydride,  176. 
Anhydrite,  448,  459. 

—  soluble,  453. 
Anions,  258. 
Anode,  15,  258. 
Anthracite,  330. 
Antimonates,  526. 
Antimonic  acids,  525. 


6o8     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Antimonious  oxide,  526. 
Antimony,  520. 

—  explosive,  521. 

—  hydride,  522. 

—  nitrate,  526. 

—  ochre,  520. 

—  pentachloride,  523. 

—  pentafluoride,  523. 
—  pentasulphide,  527. 

—  pentoxide,  525. 

—  sulphate,  526. 

—  tetroxide,  525. 

—  tribromide,  524. 

—  trichloride,  523. 

—  triiodide,  524. 

—  trioxide,  525. 

—  trisulphide,  526. 
Antimonyl  potassium  tartrate,  525. 
Apatite,  453. 

Apollinaris  water,  60. 
Aqua  regia,  225,  235. 
Aquafortis,  222. 
Aquamarine,  460. 
Aragonite,  450. 
Arc,  electric,  236,  332. 
Argentite,  436. 
Argon,  208. 

—  discovery  of,  207. 
Argyrodite,  493. 
Arsenates,  518. 
Arsenic,  511. 

—  acids,  518. 

—  disulphide,  518. 

—  oxy chloride,  516. 

—  pentasulphide,  519,  520. 
-  pentoxide,  517. 

—  tribromide,  516. 

—  trichloride,  515. 

—  trifluoride,  515. 

—  triiodide,  516. 

—  trioxide,  516. 

—  trisulphide,  371,  519. 

—  white,  516. 
Arsenious  oxide,  516. 
Arsenites,  517. 
Arseniuretted  hydrogen,  513. 
Arsenolite,  511. 

Arsine,  513. 
Asbestos,  462. 
Association,  153. 
Atmolysis,  48. 
Atmosphere,  203. 

—  and  combustion,  n. 
Atmospheric    nitrogen,    utilization    of, 

206,  236. 

—  pressure,  41. 
Atom,  definition  of,  107. 

Atomic  heat,  119.  * 


Atomic  theory,  105. 

—  volumes,  383. 

—  weights,  determination  of,  115,288, 

—  —  and  specific  heats,  118. 

—  standard  of,  1 15. 
Attraction,  molecular,  49. 
Aurates,  445. 

Auric  chloride,  446. 

—  salts,  445. 
Aurous  salts,  445. 
Avogadro's  hypothesis,  106. 
Azoimide,  219. 
Azotobakter,  206. 
Azurite,  423. 

BALANCE,  8. 
Barium,  456. 

—  chloride,  458. 

—  chromate,  538. 

—  hydroxide,  457. 

—  hypophosphite,  249. 

—  nitrate,  458. 

—  oxide,  23,  457. 

—  peroxide,  23,  458. 

—  sulphate,  458. 
Bases,  33,  98,  277. 

—  strength  of,  263,  266,  387, 
Basic  oxides,  100,  273,  371. 

—  salts,  464. 

—  slag,  555. 

Basicity  of  acids,  187,  254, 
Bauxite,  479. 
Bell  metal,  427. 
Beryllium,  460. 

—  atomic  weight  of,  385. 
Bessemer  process,  554.' 
Bismuth,  527. 

—  glance,  527. 

—  hydroxide,  530. 

—  ochre,  527. 

—  oxychloride,  529. 

—  oxynitrate,  530. 

—  peroxides  of,  530. 

—  subnitrate,  530. 

—  sulphate,  530. 

—  tribromide,  529. 

—  trichloride,  528. 

—  trifluoride,  528. 

—  triiodide,  529. 

—  trinitrate,  530. 

—  trioxide,  529. 

—  trisulphide,  530. 

"  Bismuthic  acid,"  530. 
Bismuthite,  527. 
Bismuthous  oxide,  529. 
Black  ash,  403. 
Blackband,  551. 
Blacklead,  326. 


INDEX 


609 


Blast-furnace,  552. 
Bleaching,  90,  301. 
Bleaching  powder,  452. 
Blue  vitriol,  430. 
Boiling-point  of  a  liquid,  64. 

elevation  of,  196. 

Bone- ash,  239,  453. 
Bones,  composition  of,  239. 
Boracite,  372. 
Borates,  377. 
Borax,  372,  377. 
Boric  acids,  375-377- 
Boron,  372. 

—  hydrides,  373. 

—  nitride,  374. 

—  sulphide,  375. 

—  trichloride,  374. 

—  trifluoride,  374. 

—  trioxide,  375. 
Bort,  325. 
Boyle's  law,  40. 

—  —  accuracy  of,  42,  50. 
and  kinetic  theory,  49. 

Brass,  427. 
Braunite,  543. 
Bricks,  486. 

Brin's  oxygen  process,  23, 
Britannia  metal,  522. 
Bromates,  184. 
Bromic  acid,  183. 
Bromine,  154. 

—  trifluoride,  190.     * 
-  water,  155. 

Bronze,  427. 
Brookite,  492. 
Bunsen  burner,  357. 
Burnt  alum,  485. 
Butter  of  antimony,  523. 

Cadmium,  469. 

—  compounds  of,  470. 
Caesium,  418. 
Calamine,  465. 
Calaverite,  443. 
Gale-spar,  450. 
Calcite,  450. 
Calcium,  448. 

—  bicarbonate,  340,  451. 

—  carbide,  348,  454. 

—  carbonate,  340,  350. 

—  chlorate,  413. 

—  chloride,  451. 

—  chromate,  538. 

—  cyanamide,  236,  454. 

—  fluoride,  152. 

—  hydride,  449. 

—  manganite,  545. 

—  oxide,  449. 

39 


Calcium  peroxide,  450. 

—  phosphate,  453. 

—  silicates,  454. 

—  sulphate,  452. 

—  —  solubility  in  water,  453. 

—  sulphide,  403,  405,  454. 
Calculations,  123. 
Calomel,  473. 

Caloric,  59. 
Calorific  power,  349. 
Calx,  29. 

Candle,  burning  of,  9. 
Carbamide,  343. 
Carbides,  330,  454. 
Carbon,  324. 

—  allotropic  modifications,  324. 

—  cycle,  351. 

—  dioxide,  205,  335. 
solid,  337. 

—  disulphide,  341. 

—  monoxide,  333. 

—  oxysulphide,  342. 

—  suboxide,  332. 

—  tetrachloride,  509. 
Carbonado,  325. 
Carbonates,  340. 
Carbonic  acid,  338. 
Carbonyl  chloride,  342. 
Carbonyls,  335. 
Carborundum,  330,  367. 
Carnallite,  410. 
Caro's  acid,  314. 
Cassiterite,  494. 

Cast  iron,  553. 
Catalysis,  22,  87,  140,  303. 
Cathode,  15,  258. 
Cations,  258. 
Caustic  potash,  411. 

—  soda,  398. 
Celestine,  455. 
Cement,  450. 

—  Portland,  450. 
Cementation  process,  555. 
Cementite,  557. 

Cerite,  493. 
Cerium,  493. 
Cerussite,  500. 
Chalcocite,  423. 
Chalcopyrite,  423. 
Chalk,  450. 
Chalybeate  waters,  60. 
Chamber  crystals,  305. 
Chance's  process,  405. 
Charcoal,  327. 
Charles's  law,  42. 
Chemical  affinity,  15,  272. 

—  change,  3. 

—  —  characteristics  of,  16, 


610     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Chemical  affinity,  types  of,  6. 

—  energy,  n,  144,  352. 

—  equilibrium,  164. 

in  electrolytes,  261,  439. 

Chili  saltpetre,  158,  223,  407. 
Chlorates,  181. 
Chloric  acid,  181. 
Chlorinated  lime,  452. 
Chlorine,  86. 

—  bleaching  action  of,  90. 

—  chemical  relations,  190. 

—  dioxide,  178. 

—  heptoxide,  179. 

—  hydrate,  91,  112. 

—  liquefaction  of,  71. 

—  monoxide,  177. 

—  oxides  of,  100,  178. 

—  preparation,  86. 

—  water,  89. 
Chlorites,  181. 
Chloroplatinates,  576. 
Chloroplatinites,  575. 
Chlorosulphonic  acid,  314,  316. 
Chlorous  acid,  181. 
Chromates,  538. 

Chrome  alums,  535. 

—  iron  ore,  533. 
Chromic  acids,  539. 

—  anhydride,  538. 

—  chloride,  535. 

—  hydroxide,  534. 

—  oxide,  535. 

—  sulphate,  535. 
Chromites,  536. 
Chromium,  533. 

—  ammonia  compounds,  539. 

—  trioxide,  536,  538. 
Chromous  compounds,  534. 
Chromyl  chloride,  539. 
Chrysoberyl,  460,  483. 
Cinnabar,  470. 

Classification  of  the  elements,  378. 
Clay,  370,  485- 

—  ironstone,  551. 
Coal,  330. 
Coal-gas,  348. 
Cobalt,  564. 

—  glance,  564. 
Cobaltammines,  567. 
Cobaltic  hydroxide,  565. 

—  oxide,  565. 

—  sulphate,  366. 
Cobalticyanides,  567. 
Cobaltinitrites,  567. 
Cobalto-cobaltic  oxide,  565. 
Cobaltocyanides,  567. 
Cobaltous  salts,  565. 
Coke,  328. 


Colemanite,  372. 
Collection  of  gases,  21,  39. 
Colloidal  solutions,  369,  370. 
Colloids,  reversible,  372. 
Columbite,  51*. 
Columlrium,  511. 
Combination,  chemical,  6. 
Combining  weights,  103,  115. 

law  of,  f 02. 

and  atomic  weights,  115,  120. 

chemical  equivalents,  120. 

Combustion,  26,  353. 

—  definition  of,  26. 

—  heat  of,  145. 

Complex  ions,  430,  440,  446,  578. 

—  salts,  282. 

Compound,  chemical,  4,  7. 
Concentration,  definition  of,  65. 

—  molecular,  166,  167. 
Condenser,  Liebig's,  61. 
Conductivity,  electrical,  273. 
Condy's  fluid,  550. 
Conservation  of  energy,  n. 

mass,  8. 

—  —  weight,  8 

Constant  composition,  law  of,  57,  101. 
Contact  process,  304. 
Co-ordination  numbers,  578. 
Copper,  423. 

—  alloys  of,  427. 

—  ammonia  compounds,  430. 

—  compounds.     See   under  Cuprous 

and  Cupric. 

—  equivalent  of,  126. 

—  hydrogen  arsenite,  517. 

—  properties  of,  426. 
Corrosive  sublimate,  474. 
Corundum,  479,  482. 
Critical  data,  71. 

—  phenomena,  70. 
Crocoisite,  500,  533. 
Crookesite,  488. 
Cryohydrates,  199. 
Cryolite,  149,  395,  479. 
Crystallization,  velocity  of,  293. 

—  water  of,  112,  400. 
Crystallography,  317. 
Crystalloids,  370. 
Crystals,  317. 

—  mixed,  199. 

—  symmetry  of,  317. 
Cupellation,  436. 
Cupric  carbonates,  431. 

—  chloride,  429. 

—  hydroxide.  429. 

—  nitrate,  431. 

—  oxide,  55, 

—  sulphate, 


INDEX 


611 


Cupric  sulphide,  431. 
Cuprite,  423. 
Cuprous  chloride,  428. 

—  cyanide,  428. 

—  iodide,  428. 

—  oxide,  427. 

—  sulphide,  429. 
Cyanates,  344. 

Cyanide  process  (gold),  444. 
Cyanides,  commercial  preparation,  237. 
Cyanogen,  343, 

D ALTON'S  law,  77. 
Davy  lamp,  362. 
Deacon's  process,  88. 
Decomposition,  double,  7. 

—  simple,  7. 

Degree  of  dissociation,  171,  264,  267. 
Dehydration,  402. 
Deliquescence,  403. 

Density  of  gases  and  vapours,  109,  127. 
Dewar  flasks,  75. 
Dialysed  iron,  559. 
Dialysis,  369. 
Diamond,  324. 
Dichromates,  537. 
Dichromic  acid,  539. 
Diffusion  of  gases,  46. 
law  of,  47. 

—  of  liquids,  194. 
Dimorphism,  292. 
Disintegration  of  atoms,  590. 
Dissociation,  degree  of,  171,  264. 

—  electrolytic,  261. 
and  valency,  271. 

—  thermal,  169. 
Distillation,  61. 

—  destructive  (dry),  214,  327. 

—  under  reduced  pressure,  224. 
Distribution  coefficient,  160. 
Disulphur  trioxide,  314. 
Dithionic  acid,  315. 
Dolomite,  461. 

Double  decomposition,  7. 

—  salts,  282. 

Dulong  and  Petit's  law,  118. 
Dutch  metal,  90. 
Dynamite,  225. 

EARTHENWARE,  486. 

Earth's  crust,  composition  of,  19. 

Eau  de  Javelle,  180. 

Efflorescence,  403. 

Electric  arc,  236,  348. 

—  battery,  433. 

—  furnace,  331. 
Electrical  conductivity,  257. 
Electro-affinity,  388. 


Electrochemical  equivalents,  259. 
Electrodes,  15,  258. 
Electrolysis,  15,  257. 

—  of  sodium  chloride,  398. 

—  of  water,  14. 
Electrolytes,  257. 
Electromotive  force,  432. 
Electrons,  581. 
Elements,  7,  17. 

—  classification  of,  378. 

—  list  of,  18,  and  back  of  cover. 

—  natural  occurrence,  19. 

—  potential  series  of,  434. 
Emerald,  460. 

Emery,  479,  482. 
Empirical  formula,  121. 
Endosmosis,  193. 
Endothermic  actions,  145,  172. 

—  compounds,  145,  221,  232,  342. 
Energy,  chemical,  13,  144,  352. 

—  conservation  of,  n. 

—  definition  of,  12. 

—  electrical,  13,  434. 

—  transformations  of,  12. 
Epsom  salts,  464. 
Equations,  22. 

—  writing  of,  122,  226,  537. 
Equilibria,  displacement  of,  171. 
Equilibrium,  chemical,  164. 

—  diagram  for  water,  68. 

—  displacement  of,  169,  171,  232. 

—  physical,  63  etseq. 
Equivalents,  chemical,  104,  120,    125, 

*3*« 

—  electrochemical,  259. 

—  of  acids  and  bases,  283. 
Ethane,  346. 

Ethyl  alcohol,  349. 
Ethylene,  347. 
Eutectic  mixture,  198. 
Euxenite,  493. 
Exosmosis,  193. 
Exothermic  actions,  145. 

—  compounds,  215,  338. 
Expansion  of  gases,  44. 

FACT,  113. 
Faraday's  laws,  275. 
Felspar,  4,  479. 
Ferrates,  560. 
Ferric  acid,  560. 

—  alum,  562. 

—  chloride,  561. 

—  hydroxide,  559,  561. 

—  oxide,  559. 

—  sulphate,  562. 

—  sulphide,  562. 
Ferricyanides,  563. 


6i2     A  TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Ferrocyanides,  563. 

Ferrous  ammonium  sulphate,  561. 

—  carbonate,  561. 

—  chloride,  560. 

—  hydroxide,  559. 

—  oxide,  559. 

—  sulphate,  560. 

—  sulphide,  294,  562. 
Filtration,  5. 
Fire-damp,  345. 
Flame,  354. 

—  oxy-hydrogen,  359. 

—  separator,  358. 

—  structure  of,  356. 
Flames,  luminosity  of,  355. 

—  temperature  of,  359. 
Flint,  368. 

Flowers  of  sulphur,  290. 
Fluoborates,  374. 
Fluorine,  149. 
Fluorspar,  149,  448. 
Fluosilicates,  366. 
Flux,  237. 

Formation,  heat  of,  145. 
Formulae,  empirical,  121. 

—  graphic,  131. 

—  molecular,  in,  121. 

—  significance  of,  112. 
Franklinite,  465. 

Free  affinity,  272. 
Freezing-mixture,  452. 
Freezing-point,  depression  of,  196. 
Fuming  nitric  acid,  226. 

—  sulphuric  acid,  312. 
Furnace,  blast-,  552. 

—  electric,  331. 

—  reverberatory ,  404. 
Fusible  metals,  528. 
Fusion,  heat  of,  67. 

GALENA,  500. 
Gallium,  487. 
Galvanic  battery,  433. 
Galvanized  iron,  466. 
Garnierite,  568. 
Gas  carbon,  328. 

—  laws,  40. 

deviations  from,  42,  44,  49. 

—  perfect,  50. 
Gaseous  diffusion,  46. 

—  law  of-  47. 

—  volumes,  law  of,  104. 
Gases,  collection  of,  21,  39. 

—  diffusion  of,  46. 

—  effusion  of,  48. 

—  expansion  coefficient,  44. 

—  general  equation  for,  45. 

—  general  properties  of,  40. 


Gases,  inactive,  207. 
—  kinetic  theory  of,  48. 

—  liquefaction  of,  70. 
Gay-Lussac  tower,  308. 
Gay-Lussac's  law,  104. 
Generalization,  113. 
German  silver,  427. 
Germanium,  493. 
Glass,  454. 

—  etching  of,  153. 
Glauber's  salt,  407. 
Glover  tower,  308. 
Glucinum  (see  beryllium). 
Gold,  443. 

—  alloys,  445. 

—  properties  of,  444. 
Goldschmidt's  thermite  process,  482. 
Graham's  law,  47. 
Gram-molecule,  volume  occupied  by, 

no. 

Granite,  3,  479. 
Graphic  formulae,  131. 
Graphite,  326. 
Graphitic  acid,  327. 
Gravimetric  analysis,  287. 
Greenockite,  469. 
Guncotton,  225. 
Gun-rnetal,  427. 
Gunpowder,  416. 
Gypsum,  448,  452. 

HAEMATITE,  551. 
Halogen  acids,  92,  149. 

—  —  heats  of  formation  of,  191. 
Halogens,  comparison,  190. 

—  valency  of,  341. 
Hampson's  apparatus,  72. 
Hardness  of  water,  341,  451. 
Harrogate  water,  60. 
Hartshorn,  spirits  of,  213. 
Hausmannite,  543. 

Heat,  atomic,  119. 

—  molecular,  of  gases,  208. 

—  of  combustion,  145. 

formation,  145. 

fusion,  67. 

neutralization,  269. 

solution,  145. 

vaporization,  66. 

Heavy  spar,  456. 
Helium,  209,  588. 

—  discovery  of,  207. 

—  liquefaction  of,  210. 
Henry's  law,  78. 
Hess's  law,  146. 
Heterogeneous  equilibria,  171. 
Hofmann's    method    for    determining 

vapour  densities,  129. 


INDEX 


613 


Horn  silver,  420. 
Hydrates,  113,  400. 

—  solubility  of,  400. 

—  vapour  pressure  of,  402. 
Hydrazine,  218.  » 
Hydrazoic  acid,  219. 

Hydriodic  acid,  161, 
Hydrobromic  acid,  156. 
Hydrocarbons,  332,  345. 
Hydrochloric  acid  92. 
Hydrocyanic  acid,  344. 
Hydrofluoboric  acid,  374. 
Hydrofluoric  acid,  152. 
Hydrofluosilicic  acid,  366. 
Hydrogel,  371. 
Hydrogen,  31. 

—  chemical  properties  of,  37. 

—  combination  with  oxygen,  37. 

—  occlusion  of,  39. 

—  physical  properties  of,  36. 

—  preparation  of,  31. 

—  antimonide,  522. 

—  arsenide,  513. 

—  bromide,  156. 

—  chloride,  92. 

composition  of,  96. 

—  disulphide,  296. 

—  fluoride,  152. 

—  iodide,  161. 

thermal    dissociation   of,    163, 

164,  1 66. 

—  pentasulphide,  297. 

—  peroxide,  138. 

estimation  of,  142. 

tests  for,  142. 

—  phosphides,  242,  244. 

—  poly  sulphides,  296. 

—  selenide,  320. 

—  silicide,  365. 

—  sulphide,  293. 

properties  of,  294-296. 

—  trisulphide,  297. 

—  telluride,  321. 
Hydrolysis,  252,  267,  285. 
Hydromagnesite,  464. 
Hydrosol,  371. 
Hydroxides,  254. 
Hydroxylamine,  219. 

—  salts  of,  220. 
Hypobromous  acid,  183. 
Hypochlorites.  180. 
Hypochlorous  acid,  179. 

—  anhydride,  177. 
Hypoiodous  acid,  184. 
Hyponitrous  acid,  234. 
Hypophosphoric  acid,  2^1. 
Hypophosphorous  acid,  249. 
Hyposulphurous  acid,  314. 


Hypothesis,  definition  of,  113. 

ICE,  melting-point  of,  68. 
Iceland  spar,  450. 
Ignition  temperature,  360. 
Incandescence,  355. 
Incandescent  mantles,  356. 
Indium,  487. 
Induction,  114. 
lodates,  185. 
lodic  acid,  185. 
—  anhydride,  184. 
Iodine,  158. 

—  bromide,  190. 

—  chlorides  of,  189. 

—  dioxide,  184. 

—  dissociation  of,  159,  170. 

—  pentafluoride,  190. 

—  pentoxide,  184. 
Ionic  equilibrium,  261. 
lonisation  theory,  261. 

—  degree  of,  264,  267. 
Ions,  258. 

—  complex,  440,  446,  578. 
Iridium,  573. 

Iron,  551. 

—  allotropic  modifications,  556. 

—  alloys,  554,  558. 

—  alum,  562. 

—  carbonyls,  335. 

—  cast,  553. 

—  compounds  of.    See  under  Ferrous 

and  Ferric. 

—  disulphide,  562. 

—  galvanized,  466. 

—  oxides  of,  559. 

—  passivity  of,  558. 

—  pyrites,  551,  562. 

—  rusting  of,  557. 

—  wrought,  554. 
Isodimorphism,  319. 
Isomerism,  175. 
Isomorphism,  119,  319. 
Isotopic  elements,  593. 

JASPER,  368. 
Joule-Thomson  effect,  72, 

KAINITK,  410,  461. 
Kaolin,  485. 
Kassner's  process,  504. 
Kelp,  158. 
Kieselguhr,  398. 
Kieserite,  461. 
Kinetic  theory  of  gases,  48. 
Kipp's  apparatus,  35. 
Krypton,  210. 
Kupfernickel,  512,  568. 


6i4     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


LABILE  state ,  85. 
Lakes,  483. 
Lamp-black,  328. 
Landolt,  10. 
Lanthanum,  490. 
Lapis  lazuli,  486. 
Laughing-gas,  233. 
Lavoisier,  10,  17,  27. 
Law,  definition  of,  114. 

—  Boyle's,  40. 

—  Charles's '42. 

—  Dalton's,77. 

—  Dulongand  Petit's,  118. 

—  Gay-Lussac's,  104. 

—  Graham's,  47. 

—  Henry's,  78. 

—  Hess's,  146. 

—  Newlands',  379. 

—  Periodic,  382. 

—  of  combining  weights,  103. 

—  of  constant  composition,  101. 
mass  action,  165. 

multiple  proportions,  101. 

octaves,  379. 

reciprocal  proportions,  103. 

thermoneutrality,  270. 

Laws,  Faraday's,  259. 
Lead,  500. 

—  alloys,  502. 

—  accumulator,  507. 

—  acetate,  507. 

—  bromide,  505. 

—  carbonate,  505. 

—  chloride,  504. 

—  chromate,  538. 

—  disulphate,  507. 

—  hydroxide,  503. 

—  iodide,  505. 

—  monoxide,  503. 

—  nitrate,  228,  505. 

—  peroxide,  504. 

—  properties  of,  501. 

—  red,  503. 

—  sesquioxide,  503. 

—  suboxide,  502. 

—  sulphide,  506. 

—  sulphate,  505. 

—  tetracetate,  507. 

—  tetra chloride,  507. 

—  white,  506. 

Lc  Chatelier's  theorem,  173. 
Leblanc  process,  403. 
Lepidolite,  417 
Liebig's  condenser,  61 
Lime,  449. 
Lime-water,  449. 
Limestone,  450. 
Linde-Hampscn  apparatus,  72. 


Liquation,  494. 
Liquefaction  of  gases,  70. 
Liquid  air,  73,  209,  211. 

fractionation  of,  74. 

Liquids,  diffusion  of,  194. 

—  general  properties,  62. 

—  miscibility  of,  80. 

—  vapour  pressures  of,  63. 
Lithium,  417. 

—  mica,  417. 

—  salts  of,  418. 
Lodestone,  559. 
Luminescence,  355. 
Luminosity  of  flames,  355. 

MAGNALIUM,  481. 
Magnesia,  calcined,  462. 
Magnesite,  461. 
Magnesium,  461. 

—  ammonium  arsenate,  518. 
phosphate,  464. 

—  bicarbonate,  465. 

—  carbonates,  464. 

—  chloride,  463. 

—  hydroxide,  462. 

—  nitride,  208. 

—  oxide,  462. 

—  pyrophosphate,  464. 

—  silicide,  364,  365. 

—  sulphate,  464. 
Magnetic  iron  ore,  551,  559. 
Magnetite,  559. 
Malachite,  423,  431. 
Manganates,  547. 
Manganese,  543. 

—  bronze,  427. 

—  dioxide,  88,  544. 

—  heptoxide,  545. 

—  trioxide,  545. 
Manganic  acid,  547. 

—  chloride,  546. 

—  hydroxide,  546. 

—  oxide,  544. 
Manganite,  543. 
Manganous  acid,  545. 

—  chloride,  546. 

—  hydroxide,  544. 

—  oxide,  544. 

—  sulphate,  546. 

—  sulphide,  546. 
Marble,  450. 
Marsh-gas,  345. 
Marsh's  test,  413. 

Mass  action ,  law  of,  165. 

—  active,  167. 

—  conservation  of,  8. 
Matches,  242. 
Matte,  424. 


INDEX 


Matter,  3. 

Mechanical  mixture,  4. 
Meerschaum,  461. 
Melting-point,  definition,  68. 
Mercuric  ammonium  chloride,  476. 

—  chloride,  474. 

—  cyanide,  475. 

—  diammonium  chloride,  477. 

—  iodide,  475, 

—  nitrate,  475. 

—  oxide,  474. 

—  sulphide,  476. 
Mercurous  chloride,  473. 

—  chromate,  536. 

—  iodide,  473. 

—  nitrate,  474. 

—  oxide,  473. 

—  sulphate,  474. 
Mercury,  470. 

Metals  and  non-metals,  18,  387. 

—  general  properties  of,  387. 

—  preparation  of,  390. 

—  relative     displacing     powers     of, 

434- 

Metantimonic  acid,  525. 
Metaphosphoric  acid,  253. 
Metaphosphorous  acid.  251. 
Metarsenic  acid,  518. 
Metasilicic  acid,  369. 
Metastable  state,  69,  85. 
Metastannic  acid,  498. 
Meteoric  iron,  551. 
Methane,  345. 
Microcosmic  salt,  253, 
Milk  of  lime,  449. 
Millon's  base,  477. 
Mineral  waters,  60. 
Miscibility  of  liquids,  80. 
Mixed  crystals,  199. 
Mixture,  mechanical,  4. 
Molecular  attraction,  49. 
"  Molecular  compounds,"  272,  577. 
Molecular  concentration,  167. 

—  formulae,  in,  121. 
how  to  establish,  121. 

—  —  significance  of,  112. 

—  heat  of  gases,  208. 

—  weights,  109. 

—  —  in  solution,  195. 
Molecule,  definition  of,  107. 
Molecules,  complexity  of,  108. 
Molybdenite,  540. 
Molybdenum,  540. 
Molybdic  acid,  540. 
Monoperphosphoric  acid,  254. 
Monopersulphuric  acid,  314. 
Monazite,  209,  493. 
Mordants,  483. 


Mortar,  450. 

Multiple  proportions,  law  of,  101. 

NASCENT  hydrogen,  214,  219,  514. 

—  oxygen,  91. 

—  state,  514. 
Neon,  210. 

Nessler's  reagent,  475,  477. 
Neutralization,  99. 

—  heat  of,  269. 
Newlands'  law  of  octaves,  379. 
Niccolite,  568. 

Nickel,  568. 

—  alloys  of,  568, 

—  carbonyl,  335. 

—  cyanide,  570. 

—  sulphides  of,  570. 
Nickelic  oxide,  569. 

—  hydroxide,  569.  . 
Nickelo-nickelic  oxide,  569. 
Nickelous  chloride,  570. 

—  hydroxide,  569. 

—  oxide,  569. 

—  sulphate,  570. 
Niobium,  511. 
Nitragin,  206. 
Nitramide,  235. 
Nitrates,  222,  225. 

—  "  brown  ring  test  "  for,  232. 
Nitre,  222,  394,  415. 

Nitric  acid,  222. 
fuming,  226. 

—  oxide,  231. 
Nitrides,  202,  217. 
Nitrites,  230,  407. 
Nitrogen,  200. 

—  atmospheric,  utilization  of,  206,236 
r-  utilization  by  plants,  206. 

—  bromide,  221. 

—  chloride,  221. 

—  cycle,  237. 

—  iodide,  221. 

—  pentoxide,  227. 

—  peroxide,  227. 

dissociation  of,  228. 

—  tri oxide,  229. 
Nitroglycerine,  225. 
Nitrolim,  236. 
Nitrosyl  bromide,  235. 

—  chloride,  235. 

—  fluoride,  235. 

—  sulphuric  acid,  305. 
Nitrous  acid,  230. 

—  oxide,  233. 
Nitryl  fluoride,  235. 
Noble  metals,  235. 
Nomenclature,  acids  and  salts,  188. 
Non-electrolytes,  257. 


<)i6      A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Non-metals,  18,  387. 
Nordhausen  sulphuric  acid,  312. 
Normal  solutions,  286. 

OCCLUSION,  39,  329,  573. 

Octaves,  law  of,  379. 

Oil  of  vitriol,  165,  303,  316. 

Olivine,  370. 

Opal,  368. 

Organic  acids,  350. 

—  compounds,  331. 
Orpiment,  519. 
Orthoantimonic  acid,  525. 
Orthoboric  acid,  375. 
Orthoclase,  370. 
Orthophosphoric  acid,  267. 
Orthoplumbic  acid,  504. 
Orthosilicic  acid,  370. 
Orthostannic  acid,  498. 
Osmic  acid,  573. 
Osmium,  573. 

Osmosis,  192. 
Osmotic  pressure,  192. 

and  gas  pressure,  195. 

Oxidation,  39,  433. 

Oxides,  classification  of,  273. 

Oxygen,  20. 

—  chemical  properties,  25. 

—  commercial    preparation,    23,  74, 

.5°4- 

—  discovery  of,  20. 

—  preparation,  21,  548. 

—  standard  for  atomic  weights,  115. 
Oxy-hydrogen  flame,  359. 

Ozone,  61,  133,  240. 

—  tests  for,  137. 

Ozone-oxygen   equilibrium,    135,    172, 
174. 

Palladium,  573. 

—  occlusion  of  hydrogen  by,  39,  573. 
Parkes  process,  436. 

Partial  pressures,  law  of,  77. 
Passivity  of  metals,  533,  558. 
Pattison  process,  436. 
Pearlite,  557. 
Peat,  330. 

Pentathionic  acid,  315. 
Percarbonic  acid,  341. 
Perchloric  acid,  182. 

—  anhydride,  179. 
Perchromic  acid,  539. 
Perfect  gas,  50. 
Periodic  acid,  185. 

—  law,  382. 

—  system,  379. 

deficiencies  of,  386. 

uses  of,  383. 


Permanganates,  548. 
Permanganic  acid,  548. 

—  anhydride,  545. 
Permonosulphuric  acid,  314. 
Peroxides,  140,  188,  273. 
Perphosphoric  acid,  254, 
Persulphates,  313. 
Persulphuric  acid,  313. 

—  anhydride,  313. 
Petalite,  417. 
Petzite,  443. 
Pewter,  496. 
Phase,  69. 

Phlogiston  theory,  29. 
Phosgene,  342. 
Phosphates,  251. 
Phosphine,  242. 
Phosphomolybdates,  540. 
Phosphonium  compounds,  244. 
Phosphor  bronze,  427. 
Phosphoric  acid,  239,  251. 
Phosphorite,  238,  448,  453. 
Phosphorous  acid,  250. 

—  oxide,  248. 
Phosphorus,  238. 

—  combustion  of,  28,  240,  248. 

—  diiodide  of,  247. 

—  halogen  compounds  of,  245. 

—  Hittorf  s,  242. 

—  hydrides  of,  242. 

—  metallic,  242. 

—  oxides  of,  247. 

—  oxyacids  of,  247. 

—  oxychloride,  247. 

—  pentabromide,  247. 

—  pentachloride,  246. 

dissociation  of,  169,  246. 

—  pentafluoride,  245. 

—  pentoxide,  248. 

—  red,  241. 

—  sulphides  of,  256. 

—  tetroxide,  248. 

—  tribromide,  247. 

—  trichloride,  245. 

—  trifluoride,  245. 

—  tri-iodide,  247. 

—  trioxide,  248. 

—  yellow,  240. 
Photography,  442. 
Physical  change,  2,  3. 
Pig-iron,  553. 

"  Pink  salt,"  499. 
Pitch-blende,  541,  580. 
Plaster  of  Paris,  553. 
Platinic  chloride,  576. 

—  hydroxide,  575. 

—  oxide,  575. 

—  sulphide,  576. 


INDEX 


617 


Plantinicyanides,  576. 
Platinocyanides,  576. 
Platinous  chloride,  575. 

—  hydroxide,  575. 

—  oxide,  575. 

—  sulphide,  576. 
Platinum,  574. 

—  ammonia  compounds,  576. 

—  black,  575. 

-  ore,  572,  574. 

—  spongy,  575. 
Plumbago,  326. 
Plumbates,  504. 
Plumbites,  503. 
Polonium,  580. 
Polymeride,  153,  175. 
Polymerization,  153,  175. 
Polymorphism,  175,  292. 
Polyoxides,  273. 
Polythionic  acids,  315. 
Porcelain,  486. 
Potable  waters,  60. 
Potash  alum,  485. 

—  caustic,  411. 
Potassium,  410. 

—  aluminium  sulphate,  485. 

—  antimonates,  526. 

—  antimonyl  tartrate,  525. 

—  aurichloride,  446. 

—  bicarbonate,  414. 

—  bromide,  412. 

—  carbonate,  414. 

—  carboxide,  410. 

—  chlorate,  181,  413. 

—  chloride,  412. 

—  chromate,  538. 

—  cobalticyanide,  567. 

—  cobaltinitrite,  567. 

—  cobaltocyanide,  567. 

—  cyanide,  344. 

—  dichromate,  538. 

—  ferrate,  560. 

—  ferricyanide,  563. 

—  ferrocyanide,  563. 

—  fluoride,  150. 

—  fluosilicate,  366. 

—  hydride,  411. 

—  hydrogen  fluoride,  152. 
sulphate,  415. 

—  hydroxide,  411. 

—  hypochlorite,  179,  187. 

—  hypoiodite,  184. 

—  iodide,  412. 

—  manganate,  547. 

—  mercuric  iodide,  475. 

—  metantimonate,  526. 

—  monoxide,  411. 

—  nitrate,  415. 


Potassium  nitrite,  230. 

—  oxides  of,  411. 

—  perchlorate,  182,  414. 

—  peroxide,  411. 

—  permanganate,  22,  548. 

—  persulphate,  313. 

—  platinichloride,  417,  576. 

—  polysulphides,  417. 

—  polychromates,  538. 

—  pyroantimonate,  526. 

—  silver  cyanide,  441. 

—  sulphate,  414. 

—  sulphides,  315.  416. 

Potential     series     of     the     elements, 

434- 

Pressure,  osmotic,  192. 
Properties,  characteristic,  2. 

—  non-characteristic,  2. 
Proustite,  436. 
Prussian  blue,  564. 
Puddling,  554. 

Purple  of  Cassius,  446. 
Pyrargyrite,  436. 
Pyrites,  copper,  423. 

—  iron,  551,  562. 
Pyroantimonates,  526. 
Pyroantimonic  acid,  525. 
Pyroarsenic  acid,  518. 
Pyrolusite,  543. 
Pyromorphite,  500. 
Pyrophosphoric  acid,  252. 
Pyrosulphuric  acid,  302,  306. 
Pyrosulphates,  312. 

QUARTZ,  367. 
Quicklime,  449. 
Quicksilver,  471. 

RADICAL  theory,  279. 
Radio-activity,  587. 
Radium,  587.    ' 

—  emanation,  589. 
Rain  water,  59. 
Rare  earths,  490. 
Reaction,  velocity  of,  173. 
Reactions,  reversible,  24,  87,  164. 
Reactivity,   chemical,  and  heat  of  re- 
action, 148. 

Realgar,  518. 

Reciprocal  proportions,  law  of,  103. 

Red  lead,  503. 

Reduction,  39,  433. 

Residual  affinity,  272,  579. 

Reverberatory  furnace,  404. 

Reversible  reactions,  24,  87,  164. 

Rhodium,  573. 

River  water,  60. 

Rock  crystal,  368. 


6x8     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


Rocks,  constituents  of,  370,  485. 

Rontgen  rays,  580. 

Rose's  metal,  528. 

Rouge,  559. 

Rubidium,  418. 

Ruby,  479,  482. 

Rusting  of  iron,  557. 

Ruthenium,  572. 

Rutile,  492. 

SAFETY-LAMP,  362. 
Sal  ammoniac,  213,  419. 
Salts,  basic,  282,  464. 

—  double,  282. 

—  nature  of,  99,  277. 

—  normal,  252,  281. 

—  preparation  of,  280,  390. 
Salts,  solubility  of,  82,  392. 
Sand,  363. 

Sapphire,  479,  482. 
Saturated  compounds,  347. 

—  solutions,  77. 
Scandium,  490. 
Scheele's  green,  517. 
Schonite,  464. 
Sea-water,  60. 
Selenic  acid,  321,  444. 
Selenious  acid,  321. 
Selenite,  448,  452. 
Selenium,  319. 

—  chlorides  of,  320. 

—  dioxide,  320. 
Seltzer  water,  60. 
Semi-permeable  membranes,  193. 
Senarmontite,  520. 
Serpentine,  461. 

Siemens- Martin  process,  555. 
Silent  discharge,  133. 
Silica,  367. 
Silicates,  370. 
Silicic  acids,  368. 
Silicides,  564,  365. 
Silicoethane,  365. 
Silicomethane,  365. 
Silicon,  363. 

—  carbide,  367. 

—  chloroform,  367. 

—  dioxide,  367. 

—  fluoride,  153,  365. 

—  tetrabromide,  367. 

—  tetrachloride.  367. 

—  tetraiodide,  367. 
Silver,  435. 

—  alloys,  438. 

—  antimomde,  522. 

—  arsenite,  517. 

—  arsenate,  518. 

—  bromide,  441. 


Silver  carbonate,  443. 

—  chloride,  439. 

—  chromate,  538. 

—  colloidal,  438. 

—  cyanide,  441. 

—  fluoride,  441. 

—  glance,  436. 

—  hydroxide,  438. 

—  hyponitrite,  234. 

—  iodide,  441. 

—  metaphosphate,  253. 

—  nitrate,  442. 

—  oxide,  438. 

—  periodate,  186. 

—  peroxide,  438. 

—  plating,  441. 

—  phosphate,  253. 

—  potassium  cyanide,  441. 

—  sodium  thiosulphate,  315. 

—  sub-halides  of,  439,  442. 

—  sulphate,  442. 
Slag,  424,  553. 
Slaked  lime,  449. 
Smaltite,  564. 
Smithsonite,  465. 
Soda,  calcined,  405. 

—  caustic,  398. 
Sodamide,  397. 
Sodium,  394. 

—  acetate,  345. 

—  aluminate,  480. 

—  amalgam,  397,  472. 

—  aurichloride,  446. 

—  bicarbonate,  406. 

—  bisulphate,  404. 

—  bromide,  400. 

—  carbonate,  403. 
hydrolysis  of,  340. 

—  dichromate,  538. 

—  hydride,  397. 

—  hydrogen  carbonate,  406. 

—  hydrogen  sulphate,  404. 

—  hydroxide,  397. 

—  hyposulphite,  314. 

—  metaphosphate,  252. 

—  metastannate,  498. 

—  monoxide,  397. 

—  nitrate,  223,  407. 

—  nitrite,  407. 

—  oxides  of,  397. 

—  periodate,  186. 

—  peroxide,  397. 

—  phosphates,  251,  407. 

—  platinichloride,  576. 

—  pyroantimonate,  526. 

—  pyrophosphate,  252. 

—  silicates,  408. 

—  silver  thiosulphate,  315. 


INDEX 


619 


Sodium  stannate,  498. 

—  stannite,  496. 

—  sulphate,  407. 

electrolysis  of  solution,  258. 

solubility  of,  83,  407. 

—  sulphides,  408. 

—  sulphite,  302. 

—  tetrathionate,  315. 

—  thiosulphate,  315. 
Solder,  496. 

Solid  solutions,  199,  283. 
Solubility  curves,  82. 

—  effect  of  temperature  on,  81. 

—  of  gases  in  liquids,  77. 

—  of  liquids  in  liquids,  80. 

—  of  solids  in  liquids,  81. 

—  product,  439. 

—  tables,  392. 
Solutions,  definition,  76. 

—  heat  of,  145. 

—  pressure,  85. 

Solutions,  boiling-points  of,  198. 
'—  freezing-points  of,  197. 

—  general  properties  of,  76. 

—  saturated,  77. 

—  solid,  199,  283. 

—  supersaturated,  84. 

—  unsaturated,  84. 
Solvay  process,  405. 
Soot,  328. 

Spathic  iron  ore,  551. 
Specific  heat  of  gases,  208. 

of  solids,  118. 

Spectroscope,  408. 

Spectrum  analysis,  408. 

Specular  iron  ore,  551. 

Speed  of  reaction,  173. 

Spiegeleisen,  555. 

Spinelle,  483. 

Spirits  of  hartshorn,  213. 

Spitting  of  silver,  437. 

Spodumene,  417. 

Spring  water,  59. 

Stable  state,  definition,  69,  292. 

Standard    temperature  and    pressure, 

no. 

Stannates,  498. 
Stannic  acid,  498.    . 

—  chloride,  498. 

—  hydroxide,  498. 

—  oxide,  497. 

—  sulphide,  499. 
Stannous  chloride,  496. 

—  oxide,  496. 

hydrate  of,  496. 

—  sulphide,  499. 
Stassfurt  deposits,  399,  410. 
Steam,  action  of,  on  iron,  34. 


Steam,  action  of,  on  magnesium,  33. 

—  volumetric  composition  of,  54. 
Steel,  554. 

Stibine,  522. 
Stibnite,  520. 
Strength  of  acids  and  bases,  187,  263, 

265. 

Strontianite,  455. 
Strontium,  455. 

—  chloride,  456. 

—  hydroxide,  456. 

—  nitrate,  456. 

—  oxide,  456. 

—  sulphate,  456. 
Structural  formulae,  131. 
Sublimation,  68,  337,  420. 
Substance,  4. 
Substitution,  346. 
Sulphamide,  316. 
Sulphaminic  acid,  316. 
Sulphates,  312,  391. 
Sulphides,  295. 
Sulphites,  302. 
Sulphur,  289. 

—  allotropic  modifications,  291. 

—  chlorides  of,  297, 

—  colloidal,  292. 

—  dioxide,  298. 

—  dissociation  of  vapour,  291. 

—  flowers  of,  290. 

—  heptoxide,  313. 

—  hexafluoride,  297. 

—  milk  of,  292. 

—  monoclinic,  291. 

—  nacreous,  292. 

—  plastic,  292. 

—  prismatic,  291. 

—  rhombic,  291. 

—  roll,  291. 

—  trioxide,  302,  312 
Sulphuric  acid,  333. 

—  —  chemical  properties,  309. 

—  —  fuming,  312. 

Nordhausen,  312. 

physical  properties,  309. 

preparation,  303. 

—  —  uses,  310,  311. 
Sulphurous  acid,  302. 

—  anhydride,  298. 
Sulphuryl  chloride,  316. 
Supercooling,  68,  70,  293. 
Superphosphate  of  lime,  454. 
Supersaturated  solutions,  84. 
Sylvanite  443. 

Sylvine,  410. 
Symbols,  in. 

—  list  of,  back  page  of  cover. 
Sympathetic  ink,  566. 


620     A   TEXT-BOOK   OF   INORGANIC   CHEMISTRY 


TACHYDRITE,  451,  461. 
Talc,  461. 
Tantalite,  511. 
Tantalum,  511. 
Tartar  emetic,  525. 
Telluric  acid,  322. 
Tellurous  acid,  322. 
Tellurium,  321. 

—  chlorides  of,  321. 

—  oxides  of,  322. 
Temperature,  critical,  71. 

—  ignition,  360. 

—  of  flames,  359. 
Tempering,  556. 
Tetrathionic  acid,  315. 
Thallic  salts,  489. 
Thallium,  487. 
Thallous  salts,  488. 
Theory,  definition,  114. 
Thermal  dissociation,  169. 
Thermite  process,  482. 
Thermochemical  equations,  144. 
Thermochemistry,  143. 
Thermoneutrality,  law  of,  270. 
Thioantimonates,  527. 
Thioantimonites,  527. 
Thioarsenates,  519. 
Thioarsenites,  519. 
Thiocarbonic  acid,  342. 
Thiocyanates,  344. 

Thionic  acids.  315. 

Thionyl  chloride,  316. 

Thiostannates,  499. 

Thiosulphates,  314. 

Thiosulphuric  acid,  314. 

Thomas  and  Gilchrist  process,  555. 

Thomsonite,  370. 

Thorianite,  493. 

Thorite,  493,  594. 

Thorium,  493. 

—  radio-activity  of,  592. 
Tin,  494. 

—  alloys  of,  496. 

—  amalgam ,  496. 

—  combustion  of,  28,  495. 
-  pest,  495. 

Tinplate,  496. 
Tinstone,  494. 
Tincal,  372. 
Titanium,  492. 
Tourmaline,  372. 
Transition  temperature,  292. 
Tridymite,  367. 
Triple  point,  67. 
Trithionic  acid,  315. 
Tungsten,  540. 
Tungstic  acid,  540. 
Type  metal,  522. 


Types,  theory  of,  279. 

ULTRAMARINE,  486. 
Unit  volume,  112. 
Unsaturated  compounds,  347. 
Unstable  state,  definition,  69,  292. 
Uranium,  541. 

—  radio-activity  of,  583. 
Urea,  343. 

VACUUM  vessels,  75. 

Valency,  129,  271,  279,  381,  577. 

—  and  Faraday's  laws,  271. 

—  and  structural  formulae,  129. 
Vanadinite,  510. 
Vanadium,  510. 

Vapour  densities,  determination  of,  127, 

—  pressure,  63. 
Vaporization,  heat  of,  66. 
Varec,  158. 

Velocity  of  crystallization,  293. 

—  of  reaction,  173. 
and  temperature,  174. 

Venetian  red,  559. 
Vermilion,  476. 
Vichy  water,  60. 
Vitriol,  blue,  430. 

—  green,  560. 

—  oil  of,  165,  303,  316. 
Voltameter,  14. 
Volume,  atomic,  383. 
Volumetric  analysis,  286. 

—  composition  of  steam,  54. 

WASHING  soda,  405. 
Water,  51. 

—  action  of  metals  on,  32. 

—  as  an  acid,  268. 

—  as  a  base,  268. 

—  catalytic  action  of,  334,  419. 

—  composition  by  volume,  52. 
by  weight,  55. 

—  decomposition  into  elements,  14,52. 

—  distillation  of,  61. 

—  hardness  of,  341,  451. 
-  gas,  333- 

—  glass,  408. 

—  ionisation  of,  267. 

—  of  crystallization,  113,  400. 

—  physical  properties  of,  57. 

—  purification  of,  61.- 

—  softening  of,  451. 
Waters,  natural,  59. 
Weight,  conservation  of,  8. 

—  increase  of,  in  chemical  changes,  9. 
Weights,    atomic,    table    of,   back    of 

cover. 
Weldon  process,  545. 


INDEX 


621 


Werner's  theory  of  valency,  580. 
White  lead,  506. 

—  precipitate,  476. 
Witherite,  456. 
Wolfram,  540. 
Wollastonite,  370. 
Wood,     destructive      distillation     of, 

327- 

Wood  charcoal,  327. 
Wood's  metal,  528. 
Wrought  iron,  554. 
Wulfenite,  500. 
Wiirtzite,468. 

Xenon,  210. 
Ytterbium,  490. 


Yttrium,  490 

ZEOLITES,  370. 
Zero,  absolute,  43,  49. 
Zinc,  465. 

—  alloys  of,  467. 

—  blende,  465. 

—  carbonates,  468. 

—  chloride,  467. 

—  equivalent  of,  126. 

—  hydroxide,  467. 

—  oxide,  467. 

—  sulphate,  468. 

—  sulphide,  468. 
—  white,  467. 

Zircon,  493. 
Zirconium,  493. 


PRINTED  BY   MORRISON   AND  G1BJJ   LTD.,    EDINBURGH 


Revised  ana  Enlarged  Edition.    Crown  8vo.    6s. 

OUTLINES  OF 
PHYSICAL  CHEMISTRY 


BY 


GEORGE  SENTER,  D.Sc.(Lond.),  Ph.D.,  F.I.C. 

HEAD  OF  THE  CHEMISTRY  DEPARTMENT,   B1RKBECK  COLLEGE, 
LONDON 

WITH    MANY   DIAGRAMS 

THIS  book  is  designed  to  serve  as  a  general 
introduction  to  Physical  Chemistry,  and  is 
specially  adapted  to  the  needs  of  electrical  engineers, 
to  whom  an  acquaintance  with  the  general  principles 
of  this  subject  is  becoming  of  increasing  importance. 
Particular  attention  is  devoted  to  the  theory  of  solu- 
tions and  to  the  modern  developments  of  electro- 
chemistry. The  general  principles  of  the  subject 
are  illustrated  as  far  as  possible  by  numerical  ex- 
amples, and  references  are  given  to  original  papers 
and  to  other  sources  of  information,  so  that  the 
student  may  readily  obtain  fuller  details  on  any 
point  and  learn  to  make  use  of  current  literature. 
Only  an  elementary  knowledge  of  mathematics  is 
assumed. 


REVIEWS 

OF  THE  FIRST  EDITION   OF 

"  Outlines  of  Physical  Chemistry  " 
by  Dr.  Senter 

W.  OSTWALD  in  "Zeitschrift  fiir  physikalische  Cheinie" 

"The  book  not  only  belongs  to  the  class  to  which  the  description 
'  good  '  can  be  applied,  but  takes  high  rank  in  that  class.  It  is  clearly 
and  accurately  written." 

NATURE 

"  On  the  whole  the  book  is  one  which  can  be  recommended  to  all  who 
wish  to  obtain  a  first  acquaintance  with  the  subject  of  Physical  Chemistry. 
In  language  it  is  clear  and  well  expressed,  and  the  practical  illustrations 
which  are  appended  to  most  of  the  chapters  will  be  found  very  useful  for 
laboratory  work." 

AMERICAN   CHEMICAL  JOURNAL 

"...  The  various  sections  are  closed  by  excellent  practical  illustrations. 
The  presentation  is  exceptionally  clear.  The  book  is  one  of  the  best 
elementary  texts  that  have  appeared,  and  deserves  a  cordial  welcome." 

KNOWLEDGE 

"  The  hand  of  the  expert  teacher  is  manifest  on  every  page,  and  we  can 
thoroughly  recommend  the  book  as  a  clear  and  efficient  guide  to  the 
physics  of  Chemistry." 

SCIENCE   PROGRESS 

"The  author's  familiarity  with  the  student's  point  of  view  enables  him 
to  expound  with  commendable  clearness  the  leading  principles  of  this 
attractive  branch  of  science." 

ZEITSCHRIFT   FUR  ELEKTROCHEMIE 

"The  subject-matter  is  thoroughly  up  to  date,  and  the  book  itself  will 
serve  admirably  as  an  introduction  to  the  more  specialised  treatises  on  the 
various  branches." 

FARADAY  HOUSE  JOURNAL 

"  The  book,  the  text  of  which  is  illustrated  by  many  diagrams,  is  clearly 
written,  and  this,  coupled  with  the  very  moderate  price  at  which  it  is  pub- 
lished, should  do  much  to  popularise  the  study  of  Physical  Chemistry." 

SCHOOLMASTER 

"  Dr.  Senter's  experience  and  success  as  a  lecturer  in  London  enables 
him  to  treat  his  subject  on  both  the  theoretical  and  practical  sides  with  a 
thoroughness  which  makes  the  book  one  to  be  commended." 

METHUEN  &  CO.  LTD.,  ESSEX  STREET,  LONDON,  W.C 


5692 19 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


THIS  BOOK  IS  DUE  ON  THE  LAST  DATE 


P 

w 

AN  INITIAL  FI! 

WILL  BE  ASSESSED  FO 
THIS  BOOK  ON  THE  D; 
WILL  INCREASE  TO  SO 
DAY  AND  TO  $1.OO  C 
OVERDUE. 

^E  OF  25  CENTS     

R    FAILURE  TO    RETURN 
VTE   DUE.   THE  PENALTY        Vt- 
CENTS  ON  THE  FOURTH         l6' 
5N    THE    SEVENTH     DAY 

E 

0 

96 

144 

20 

58.5 
222.5 

J4 

I9^ 
16 

106.5 
3i 
195 
39 

141 
226 
103 

85.5 
101.5 

iSO.5 
44 
79 
28 
108 

23 

87.5 
32 
181.5 

127-5 
159 
204 

232-5 
168.5 
119 

48 
184 
238.5 
5* 
130 

173-5 

89 
65-5 
90.5 

decimal 
r    many 

Alii 

An 

Arc 

OCT  5  1943 

2 
68 

Ars 

Bai 

OCc- 

.n.  T»      4-A-rf-J 

01 

Boi 
Brc 
Ca< 
Caf 
Ca 

NOV   24  IS** 

!oo 

MAR  29  |84T 

.04 
.2 

Ca 
Cei 

.10 
g 

Ch 
Ch 
Co 

.0 

.9 

Co 
Co 

Dy 
Erl 
Eu 
Fh 
Ga 
Ga 
Ge 
Gl 
( 

.45 
.7 
•1 

:.l 
.2 

..3 
.88 

.0 

'.63 

:.06 
fi 

ft 

a< 

'.5 
1.2 

4 

n 

t.O 
J.4 
$.5 

r< 

c 

5.7 
>  -I 

1.0 
5.2 

( 

vl 
vl 
vl 

T 

1C 

r 

L.O 
).2 

5  5 

}.7 

5.37 
).6 

r  of 

LD  21-100w.-7,>39(402s) 

